V 


SOLUBILITIES  OF  BASES  AND  SALTS  IN  WATER  AT  18°. 


K 

Na 

Li 

Ag 

Tl 

Ba 

Sr 

Ca 

Mg 

Za 

FB 

Cl 

32.95 
3.9 

35.86 
5.42 

77.79 
13.3 

0.0313 
0.059 

0.3 
0.013 

37.24 
1.7 

51.09 
3.0 

73.19 
5.4 

55.81 
5.1 

203.9 
9.2 

1.49 
0.05 

Br 

65.86 
4.6 

88.76 
6.9 

168.7 
12.6 

0.0.1 
0.066 

0.042 
0.0215 

103.6 
2.9 

96.52 
3.4 

143.3 
5.2 

103.1 
4.6 

478.2 
9.8 

0.598 
0.02 

I 

137.5 
6.0 

177.9 
8.1 

1615 

8.5 

0.0G35 
0.071 

0.006 
0.0817 

201.4 
3.8 

169.2 
3.9 

200 
4.8 

148.2 
4.1 

419 
6.9 

0.08 
0.0a2 

F 

92.56 
12.4 

4.44 
1.06 

027 
0.11 

105.4 
13.5 

72.05 
3 

o.i  a 

0.0,92 

0.012 
0.001 

0.0016 
0.0,2 

0.0087 
0.0,14 

0.005 
0.085 

0.06 
0.002 

JNO, 

30.34 
2.6 

8397 
7.4 

7143 
7.3 

213.4 
8.4 

8.91 
0.35 

8.74 
0.33 

66.27 
2.7 

121.8 
5.2 

74.31 
4.0 

117.8 
4.7 

51.66 
1.4 

ClOa 

6.6 
0.52 

97.16 
6.4 

313.4 
15.3 

12.25 
0.6 

3.69 
0.13 

35.42 
1.1 

174.9 
4.6 

179.3 
5.3 

126.4 
4.7 

183.9 
6.3 

150.6 
3.16 

3rO8 

6.38 
0.38 

36.67 
2.2 

152.5 
8.20 

0.59 
0.025 

0.30 

0.009 

0.8 
0.02 

30.0 
0.9 

85.17 
2.3 

42.86 
1.5 

58.43 
1.8 

1.3 

0.03 

ioa 

7.62 
0.35 

8.33 
0.4 

80.43 
3.84 

0.004 
0.0314 

0.059 
0.021G 

0.05 
0.001 

0.25 
0.0,57 

0.25 
0.007 

6.87 
0.26 

0.83 
0.02 

0.002 
0.043 

OH 

142.9 
18 

116.4 
21. 

12.04 
5.0 

0.01 
0.001 

40.04 
1.76 

3.7 
0.22 

0.77 
0.063 

0.17 
0.02 

0.001 
0.082 

0.0,5 
0.045 

0.01 

0.0,4: 

S04 

11.11 
0.62 

16.83 
1.15 

35.64 
2.8 

0.55 
0.020 

4."4 
0.09 

0.0323 
0.0410 

0.011 
0.036 

0.20 
0.015 

35.43 
2.8 

53.12 
3.1 

0.0041 
0.0a13 

GrO4 

63.1 
2.7 

61.21 
3.30 

111.6 
6.5 

0.0025 
0.0315 

0.006 
0.031 

0.0,35 
0.0414 

0.12 
0.006 

0.4 
0.03 

73.0 
4.3 

0.0.2 
0.065 

C20, 

30.27 
1.6 

3.34 

0.24 

7.22 
O.G9 

0.0034 
0.031V 

148 
0.030 

0.0085 
0.0338 

0.0046 
0.0326 

0.0,55 
0.0443 

0.03 
0.0027 

0.0364 
0.044 

0.0.16 

0.0554 

C03 

108.0 
5.9 

19.39 
1.8 

1.3 
0.17 

0.003 
0.031 

4.95 
0.10 

0.0023 
0.0311 

0.0011 
0.047 

0.0013 
0.0313 

0.1 
0.01 

0.004? 
0.033? 

0.0,1 
0.043 

The  upper  number  in  each  square  gives  the  number  of  grams  of  the 
anhydrous  salt  held  in  solution  by  100  c.c.  of  water.  The  lower  number  is 
the  molar  solubility,  i.e.,  the  number  of  moles  contained  in  one  liter  of  the 
saturated  solution.  The  numbers  for  small  solubilities  have  been  abbreviated. 
Thus  0.0e4  -  0.0000004.  For  some  other  solubilities,  see  pages  191,  656. 


JOHN  ALEXANDER  JAMESON,  JR. 
1903-1934 


ENGINEERING  LIBRARY 


THIS  BOOK  belonged  to  John  Alexander  Jameson,  Jr.,  A.B.,  Wil- 
liams, 1925;  B.S.,  Massachusetts  Institute  of  Technology,  1928; 
M.SU  California,  1933.  He  was  a  member  of  Phi  Beta  Kappa,  Tau 
Beta  Pi,  the  American  Society  of  Civil  Engineers,  and  the  Sigma 
Phi  Fraternity.  His  untimely  death  cut  short  a  promising  career. 
He  was  engaged,  as  Research  Assistant  in  Mechanical  Engineering, 
upon  the  design  and  construction  of  the  U.  S.  Tidal  Model  Labora- 
tory of  the  University  of  California. 

His  genial  nature  and  unostentatious  effectiveness  were  founded 
on  integrity,  loyalty,  and  devotion.  These  qualities,  recognized  by 
everyone,  make  his  life  a  continuing  beneficence.  Memory  of  him 
will  not  fail  among  those  who  knew  him. 


BY  THE  SAME  AUTHOR 

EXPERIMENTAL 
INORGANIC  CHEMISTRY 

Fifth  Edition.  Revised 


A  briefer  text,  on  the  same  plan 
GENERAL    CHEMISTRY   FOR   COLLEGES 

Second  Edition.  Revised  (1916) 

Eighty-third  Thousand 
662  +x  pp.    With  138  Figures 


LABORATORY  OUTLINE  OF 
COLLEGE  CHEMISTRY 

206  +  x  pp.    With  30  Figures 


A  Textbook  of 
ELEMENTARY  CHEMISTRY 

439  +  v"i  PP-    With  98  Figures  and  6  Plates 


A  LABORATORY  OUTLINE  OF 
ELEMENTARY  CHEMISTRY 

137  pp.    With  1 8  Figures 


NEW  YORK.  THE  CENTURY  CO. 
LONDON,  GEO.  BELL  AND  SONS 


INTRODUCTION 

TO 

INORGANIC   CHEMISTRY 


BY 

ALEXANDER   SMITH 

PBOFESSOR  OF  CHEMISTRY  AND  ADMINISTRATIVE  HEAD  OF  THE  DEPARTMENT 
OF  CHEMISTRY  IN  COLUMBIA  UNIVERSITY,  NEW  YORK 


Ttbirfc  EMtfon 

REWRITTEN 


NEW  YORK 

THE    CENTURY    CO. 
1923 


COPYRIGHT,  1905, 1906, 1917, 

BT 

THE  CENTURY  CO. 


UfifiARY 


Printed  in  U.  S.  A. 


PREFACE   TO   THE   FIRST   EDITION 


THIS  book,  the  first  draft  of  which  was  written  six  years  ago,  is  the 
outgrowth  of  the  introductory  course  in  chemistry  which  the  author 
has  given  for  the  past  fifteen  years.  A  subject  undergoing  the  per- 
sistent, though  unconscious  criticism  of  keen  minds  should  gain  in 
self-consistency  and  coherence  as  it  is  presented  year  after  year.  For 
example,  an  answer  must  be  found  for  the  common  question,  "  Why 
does  the  chemistry  of  the  laboratory  differ  from  the  chemistry  of  the 
text-book  and  the  lecture  to  such  an  extent  that  they  seem  to  be  different 
sciences  ?  "  The  chemistry  of  the  laboratory  is,  of  course,  the  only 
real  chemistry,  and  that  of  the  lecture  must  be  somewhere  at  fault. 
The  student  neither  sees  nor  weighs  atoms,  for  instance,  and  so  the 
details  of  the  laboratory  experiment,  which  are  seen  and  studied,  become 
the  basis  of  the  whole  treatment.  The  atom  and  the  ion  assume  the 
role  of  merely  figurative  aids  in  the  description  of  the  facts.  Gradually 
the  conception  of  chemical  equilibrium  comes  to  contribute  the  major 
part  of  the  explanation  which  is  essential  to  the  evolution  of  a  system 
of  chemistry  founded  upon  experiment. 

In  the  choice  and  arrangement  of  the  material,  several  principles 
have  served  as  guides  : 

The  book  is  intended  primarily  for  students  beginning  the  study  of 
chemistry  in  a  college,  university,  or  professional  school.  It  is  assumed 
that  use  of  the  book  goes  hand  in  hand  with  systematically  arranged 
laboratory  work  in  general  chemistry.  The  first  four  chapters,  for  ex- 
ample, contain  a  discussion  of  a  few  typical  experiments.  They  appeal 
directly  to  experience  derived  from  the  performance  and  observation 
of  these  and  other  similar  experiments  in  the  laboratory  and  in  the 
class-room.  In  these  chapters  some  of  the  features  which  are  charac- 
teristic of  every  chemical  phenomenon  are  sought  out,  puf  into  words, 
and  illustrated. 


889745 


Vl  PREFACE 

No  conception  is  defined,  and  no  generalization  or  law  is  developed, 
until  such  a  point  has  been  reached  that  applications  of  the  conception 
and  experimental  illustrations,  later  to  be  related  in  the  law,  have 
already  been  encountered,  and  there  is  about  to  be  occasion  for  further 
applications  and  illustrations  of  the  same  things  in  the  chapters 
immediately  succeeding.  In  these  chapters  the  applications  are 
frequent  and  explicit.  Later,  page  references  in  parentheses  con- 
tinue to  indicate  the  recurrence  of  examples  which  might  other- 
wise fail  to  be  noticed.  It  is  one  thing  to  come  to  know  a  principle 
of  the  science,  and  quite  another  thing  to  have  acquired,  by  constant 
repetition  of  the  process,  a  confirmed  habit  of  using  the  prin- 
ciple on  every  appropriate  occasion.  To  assist  still  further  in  the 
attainment  of  this  end,  an  attempt  is  made  in  the  last  six  chapters 
again  to  mention  and  illustrate,  by  way  of  review,  the  most  important 
principles  of  the  science. 

No  conception  or  principle  is  given  at  all,  unless,  in  its  most  elemen- 
tary aspects,  it  can  be  made  clear  to  a  beginner  ;  and  unless  it  is 
capable  of  numerous  applications  in  elementary  work;  and,  finally, 
unless  a  knowledge  of  it  is  of  material  use  in  organizing  and  unifying 
the  result  of  such  elementary  work. 

An  attempt  has  been  made  to  state  the  laws  and  to  define  the  con- 
ceptions of  the  science  in  terms  of  experimental  facts.  The  figurative 
language  of  hypothesis  has  been  employed  only  in  explanations. 

Familiarity  with  physical  conceptions  and  facts  is  so  indispensable 
to  the  chemist  that  no  apology  is  needed  for  the  rather  full  treatment 
which  some  of  them  have  received. 

No  two  chemists  would  agree  perfectly  in  regard  to  the  apportion- 
ment of  space.  The  processes  of  chemical  industry,  and  the  every-day 
applications  of  chemical  science,  cannot  all  be  mentioned.  These 
fields,  and  that  of  mineralogy,  can  be  represented  by  examples,  with- 
Diit  the  incompleteness  of  the  result  being  in  any  way  a  detriment  to  a 
work  of  a  general  character.  Again,  a  dense  array  of  descriptive 
material,  unillumined  by  explanation,  is  a  positive  injury  to  an  intro- 
ductory treatise.  All  reference  to  historical  matters  cannot  be  omitted, 
but  a  logical  display  of  the  subject  can  be  achieved  with  comparatively 


PREFACE  Vll 

little  of  the  history  Of  all  the  aspects  of  the  science,  the  theoretical 
is  thus  the  one  whose  treatment  is  susceptible  of  least  abbreviation. 

The  principles  of  chemical  equilibrium  are  (and  have  been  for  the 
past  half-century)  fully  as  much  required  for  intelligent  consideration 
of  the  simplest  experiment,  as  is  the  theory  of  combining  proportions 
itself.  Important  parts  of  the  theories  of  solutions  and  of  the  battery 
are  much  more  recent,  but  each  is  equally  indispensable  to  the  under- 
standing of  matters  which  cannot  long  be  withheld  from  the  notice  of 
the  beginner.  Surely  space  ought  not  to  be  saved  by  entire  omission  of 
essential  parts  of  the  chief  thing  that  makes  chemistry  at  all  worthy 
of  a  place  amongst  the  sciences.  Nor  may  we  attain  brevity,  no  matter 
how  great  the  temptation,  by  condensing  the  passages  on  theory  until 
they  reach  the  limit  of  comprehensibility  by  an  expert.  Without  clear 
exposition,  full  illustration,  and  frequent  application,  laws  and  princi- 
ples simply  repel,  or  worse  still,  mislead  the  beginner. 

We  reach  the  same  conclusion  from  another  view-point.  Every 
student  should  have  access  to,  and  should  use,  reference  books  devoted 
especially  to  descriptive,  industrial,  historical,  and  physical  chemistry, 
and  to  mineralogy  and  crystallography  —  at  least  one  good  book  in 
each  of  these  five  subjects.  With  the  help  of  the  index,  the  veriest  tyro 
can  find  in  a  few  moments,  almost  anything  he  wants  in  four  out  of 
five  of  these  branches.  But  just  the  opposite  is  the  case  with  the 
theory.  Only  an  expert  realizes  what  information  he  is  in  need  of, 
and  knows  under  what  titles  to  look  for  it.  And  often  even  the 
expert  would  fail  to  understand  the  isolated  sentence  or  paragraph 
when  found.  In  a  large  proportion  of  connections  the  beginner  sim- 
ply cannot  use  such  a  book  for  rapid  reference  at  all.  In  many  lines, 
therefore,  much  may  be  left  to  outside  reading,  but  for  theory  almost 
no  dependence  can  be  placed  on  reference  work  in  other  books. 

For  the  reasons  enumerated  above,  an  unusually  large  proportion 
of  space  has  been  given  to  theoretical  matters.  The  actual  amount  of 
theory  is  no  greater  than  is  usual  in  books  of  the  same  class,  but  the 
explanations  are  often  fuller.  Even  so,  the  beginner  will  probably 
find  that  some  parts  form  reading  as  stiff  as  any  he  is  accustomed  to 
undertake,  without  complaint,  in  physics  or  mathematics.  It  can  only 


Vlll  PREFACE 

be  said  that  easily  read  modes  of  presenting  the  science  of  chemistry 
are  apt  to  delude  the  beginner  into  thinking  he  has  mastered  the  sub- 
ject, when  in  reality  he  has  simply  been  steered  clear  of  the  chief  diffi- 
culties. 

The  order  of  topics  was  determined  by  many  considerations, 
jointly.  For  example,  in  the  first  week  of  his  work,  a  student  may 
encounter  experiments,  in  connection  with  which,  almost  every  part  of 
chemical  theory  might  usefully  be  discussed.  But  mastery  of  the 
theory  must  necessarily  come  bit  by  bit,  and  the  theory  is  therefore 
distributed  through  the  book.  Instead  of  being  introduced  as  soon  as 
a  fragment  offers  a  chance  for  explanation,  the  treatment  of  each  of  the 
various  theoretical  subjects,  as  far  as  possible,  has  been  postponed  until 
a  whole  chapter  could  be  devoted  to  it.  The  result  makes  subsequent 
reference  easier,  and  facilitates  alterations  in  the  order  of  study. 
Thus,  the  hypothesis  of  ions  is  not  mentioned  as  soon  as  it  well  might 
be,  because  satisfactory  treatment  of  it  must  follow  the  molecular  and 
atomic  hypotheses  of  which  it  is  an  extension,  and  because  the  full 
explanation  of  this  hypothesis  must  be  preceded  by  some  account  of  the 
phenomena  of  electrolysis  and  of  the  essential  properties  of  solutions, 
and,  also,  by  a  discussion  of  chemical  equilibrium,  a  subject  which  of 
necessity  presupposes  two  or  three  months'  work  in  chemistry.  There 
is  another  disadvantage  which  arises  from  a  premature  explanation  of 
the  hypothesis  of  ionization.  When  it  appears  at  an  early  stage,  too 
long  an  interval  separates  this  subject  from  the  study  of  the  metallic 
elements,  and  the  details  are  largely  forgotten  before  the  field  for  their 
chief  employment  is  reached. 

The  paragraphs  in  smaller  type  are  not  intended  for  beginners,  but 
for  advanced  students  and  teachers,  which  accounts  for  the  fact  that 
reference  will  frequently  be  found  in  them  to  subjects  treated  system- 
atically in  later  chapters. 

The  exercises  and  problems  are  simply  samples  of  some  of  the 
various  kinds  of  questions  which  might  be  raised  in  dozens  at  the  end 
of  every  chapter. 

Recent  works  on  general  chemistry  have  been  consulted  during  the 
revision  of  the  manuscript.  Of  these  A.  A.  Noyes*  admirable  General 


PREFACE  IX 

y  Ostwald's  Grundlinien  —  a  veritable  tour  de  force,  and  Blox- 
arn's  Chemistry  may  be  mentioned  as  having  proved  most  suggestive. 
The  author  owes  special  thanks  to  several  friends  who  have  undertaken 
the  toilsome  work  of  reading  part  or  all  of  the  book  in  manuscript  or  in 
proof,  and  in  particular  to  his  colleagues  Messrs.  Julius  Stieglitz,  H.  N. 
McCoy,  L.  W.  Jones,  and  E.  S.  Hall,  to  Dr.  J.  B.  Tingle  of  Johns 
Hopkins  University,  to  Mr.  C.  M.  Wirick  of  the  R.  T.  Crane  High 
School,  Chicago,  and  to  Mr.  Maurice  Pincoffs  of  Chicago.  The  author 
alone  is  responsible  for  any  defects  which  may  be  inherent  in  the  plan 
of  the  book  and  for  errors  which  may  have  escaped  detection,  but  must 
gratefully  acknowledge  the  very  great  benefit  the  book  has  derived 
from  the  friendly  criticism  of  these  gentlemen.  Other  corrections  or 
suggestions  will  be  gladly  received  by  the  author. 

CHICAGO,  January,  1906.  ALEXANDER  SMITH. 


PREFACE   TO   THE   THIRD   EDITION 


THE  general  arrangement  of  the  book  has  not  been  altered,  ex- 
cepting that  the  difficult  chapter  on  the  oxygen  acids  of  chlorine  has 
been  transferred  to  a  later  position.  The  contents  have  been  brought 
up  to  date.  The  introductory  chapters  have  been  improved.  More 
applications  of  chemistry  have  been  introduced.  Various  methods 
of  writing  equations  are  discussed.  Additional  paragraphs  for  ad- 
vanced students  (in  small  type)  dealing  with  various  points  of  view 
in  regard  to  subjects  like  valence,  and  with  the  logical  arrangement 
of  matters  like  chemical  properties,  have  been  added.  Experiments 
suited  for  demonstration  purposes  have  been  designated  by  the  sign 
[Lect.  exp.]. 

In  order  to  induce  the  student  to  carry  the  book  to  and  from  the 
laboratory,  the  volume,  as  in  the  previous  editions,  has  been  made  as 
compact  as  possible  by  the  use  of  narrow  margins  and  specially 
made,  thin  paper. 

The  author  is  greatly  indebted  to  many  friends,  whose  suggestions 
have  been  utilized  in  the  revision.  He  is  also  under  great  obli- 
gations to  Dr.  J.  E.  Booge,  Dr.  George  Scatchard,  and  Mr.  Kenneth 
P.  Monroe  for  cooperating  in  the  reading  of  the  proofs,  and  for  the 
many  corrections  and  improvements  to  which  they  have  called  his 
attention. 

NEW  YORK,  January  1917,  ALEXANDER  SMITH. 


CONTENTS 


CHAPTER 

I.   CHEMICAL  PHENOMENA  AND  THE  METHODS  OF  STUDYING  AND 

CLASSIFYING  THEM 1 

II.   ENERGY    IN    CHEMICAL    CHANGE.      PHYSICS    IN    PRACTICAL 

CHEMISTRY 28 

III.  COMBINING  PROPORTIONS  BY  WEIGHT 52 

IV.  SYMBOLS,  FORMULAE,  EQUATIONS,  CALCULATIONS 69 

V.   OXYGEN 79 

VI.   MEASUREMENT  OF  QUANTITY  IN  GASES 103 

VII.  HYDROGEN 114 

VIII.   WATER 141 

IX.   RELATIONS   BETWEEN   THE    STRUCTURE    AND    BEHAVIOR   OF 

MATTER.     THE  KINETIC  MOLECULAR  VIEWPOINT 160 

X.   SOLUTION 178 

XL   HYDROGEN  CHLORIDE  AND  CHLORINE 206 

XII.   MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS 231 

XIII.  APPLICATIONS  OF  MOLECULAR  AND  ATOMIC  WEIGHTS.    PROP- 

ERTIES OF  ATOMS 255 

XIV.  THE  HALOGEN  FAMILY 267 

XV.   CHEMICAL  EQUILIBRIUM 287 

XVI.  OZONE  AND  HYDROGEN  PEROXIDE 311 

XVII.   DISSOCIATION  IN  SOLUTION 324 

XVIII.   IONIZATION 342 

XIX.   IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS 372 

XX.   SULPHUR  AND  HYDROGEN  SULPHIDE 410 

XXI.  THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR 424 

XXII.  SELENIUM  AND  TELLURIUM.    THE   CLASSIFICATION  OF  THE 

ELEMENTS 453 

XXIII.  OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS.    OXIDATION 

AND  REDUCTION  . 472 

XXIV.  THE  ATMOSPHERE.    THE  HELIUM  FAMILY 499 

xiii 


XIV  CONTENTS 

CHAPTER  PAGE 

XXV.   NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN 513 

XXVI.   OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN 525 

XXVII.   PHOSPHORUS 547 

XXVIII.   CARBON  AND  THE  OXIDES  OF  CARBON 566 

XXIX.   THE  HYDROCARBONS,  ILLUMINANTS,  FLAME 585 

XXX.   THE    CARBOHYDRATES,    ORGANIC    ACIDS,    ALCOHOLS,    SOAP, 

COLLOIDS,  FOODS 603 

XXXI.  SILICON  AND  BORON 630 

XXXII.   THE  BASE-FORMING  ELEMENTS 642 

XXXIII.  METALLIC   ELEMENTS   OF   THE   ALKALIES:     POTASSIUM   AND 

AMMONIUM 661 

XXXIV.  SODIUM    AND    LITHIUM.     IONIC    EQUILIBRIUM    CONSIDERED 

QUANTITATIVELY 683 

XXXV.   METALLIC    ELEMENTS    OF    THE    ALKALINE    EARTHS:    CAL- 
CIUM, STRONTIUM,  BARIUM  . 701 

XXXVI.   COPPER,  SILVER,  GOLD 734 

XXXVII.   MAGNESIUM,     ZINC,     CADMIUM,   MERCURY.     THE  RECOGNI- 
TION OF  THE  CATIONS  IN  QUALITATIVE  ANALYSIS 763 

XXXVIII.   ELECTROMOTIVE  CHEMISTRY 786 

XXXIX.   ALUMINIUM  AND  METALLIC  ELEMENTS  OF  THE  EARTHS 807 

XL.   GERMANIUM,  TIN,  LEAD 822 

XLI.   ARSENIC,  ANTIMONY,  BISMUTH 839 

XLII.   THE  CHROMIUM  FAMILY.     RADIUM 854 

XLIII.   MANGANESE 875 

XLIV.   IRON,  COBALT,  NICKEL. 884 

XLV.   THE  PLATINUM  METALS 905 

APPENDIX.  .  909 


INTRODUCTION   TO 
GENERAL    INORGANIC    CHEMISTRY 


INTRODUCTION   TO 
INORGANIC    CHEMISTRY 


CHAPTER  I 

CHEMICAL  PHENOMENA,   AND  THE  METHODS  OF   STUDYING 
AND   OF   CLASSIFYING  THEM 

HUMAN  knowledge  has  become,  in  recent  times,  so  extensive  and 
complex  that  the  truths  ascertained  have  had  to  be  divided;  more  or 
less  arbitrarily,  into  groups.  Thus,  the  study  of  animals,  their  classi- 
fication by  structure,  life-history,  distribution,  and  so  forth,  form  the 
group  known  as  zoology.  Such  a  group  is  called  a  science,  and  in- 
cludes a  more  or  less  distinct  body  of  knowledge.  There  is  a  widely 
spread  impression  that  a  science,  like  chemistry,  is  a  part  of  the 
natural  order  of  the  Universe.  It  is  thought  that  we  are  trying  to 
find  the  boundaries  of  chemistry,  as  they  have  been  predetermined 
by  nature,  and  to  discover  the  facts,  relations  of  facts,  and  laws 
which  nature  has  provided  as  a  means  of  classifying  the  content 
of  the  science.  Now,  the  situation  is  precisely  the  reverse  of  this. 
Nature  provides  only  the  materials  and  the  phenomena,  and  man  is 
attempting  to  classify  them.  He  divides  the  whole  into  groups,  such 
as  physics,  chemistry,  botany,  etc.  Then  he  classifies  the  facts 
within  each  group,  in  order  that  he  may  more  easily  remember  them 
and  perceive  their  relations.  He  often  finds  that,  when  new  facts 
are  discovered,  parts  of  the  classification  have  to  be  changed. 

That  the  boundaries  of  these  groups  are  purely  arbitrary,  how- 
ever, and  do  not  exist  in  the  subject-matter  itself,  is  seen  at  once  in 
our  own  treatment  of  them.  Thus,  for  convenience,  we  take  the 
structure  of  animals  by  itself  and  style  it  anatomy;  but  we  include 
in  the  science  of  physiology  the  study  of  the  way  in  which  the  parts 
of  both  plants  and  animals  perform  their  functions:  and  we  assemble 
cognate  parts  of  two  groups  in  sciences  like  astro-physics  and  phys- 
ical chemistry.  The  sciences,  therefore,  are  not  mutually  exclusive, 
and  their  boundaries  overlap  in  every  direction. 

1 


2  INORGANIC  CHEMISTRY 

The  difficulty  in  decidmg  what  are  the  most  convenient  boundaries 
is  as  great  with  chemistry  as  with  ether  groups.  At  the  one  ex- 
treme we  have  the  abstract  sciences,  logic  and  mathematics.  At 
the  other  extreme  lie  the  concrete  sciences  like  geology,  zoology,  and 
astronomy.  The  former  are  not  concerned  primarily  with  the  study 
of  matter  at  all,  but  with  that  of  abstract  conceptions.  The  latter 
deal  with  definite  aggregates  of  matter,  such  as  the  nature  and  his- 
tory of  a  particular  deposit  of  sulphur  and  their  relation  to  those 
of  other  deposits  of  sulphur,  or  the  structure  and  history  of  a  particu- 
lar collection  of  organic  material  known  as  a  pike,  and  their  relation 
to  the  structure  and  history  of  a  mass  of  similar  material  called  a 
salmon.  Between  these  two  sets  of  sciences  are  the  regions  occupied 
by  the  abstract-concrete  sciences,  physics  and  chemistry.  These 
sciences  deal  in  part  with  the  same  portions  of  matter,  but  in  a  more 
abstract. way  than  do  geology  or  biology.  To  them,  all  specimens  of 
pure  sulphur  are  alike,  whether  they  have  been  formed  by  volcanic 
action,  or  have  been  deposited  by  bacteria  in  an  entirely  different 
manner.  In  particular,  the  line  which  divides  physics  and  chemistry 
from  one  another  is  often  difficult  to  draw.  It  is  assumed,  however, 
that  the  reader  is  already  familiar  with  the  elements  of  physics,  and 
so,  in  place  of  entering  upon  an  academic  discussion  of  the  nature  of 
this  line,  we  shall  allow  its  location  to  emerge  as  we  proceed. 

The  same  principle  of  grouping  is  pursued  within  each  field. 
Thus  the  preparation  and  properties  of  chemical  compounds  is  called 
descriptive  chemistry.  The  content  of  this  portion  of  the  subject 
is  in  turn  divided  into  organic  chemistry  (dealing  with  almost  all 
the  compounds  containing  carbon)  and  inorganic  chemistry,  and 
the  content  of  each  of  these  is  further  classified  according  to  a  plan 
involving  the  consideration  of  the  constituents  of  each  compound. 
The  study  of  the  proportions  of  the  constituents  in  compounds, 
of  the  conditions  under  which  chemical  action  occurs,  and  related 
matters  of  a  more  abstract  character,  are  grouped  together  in  theo- 
retical chemistry,  which  is  likewise  subdivided.  Again,  the  means 
that  have  been  devised  for  recognizing  the  components  of  mixtures 
or  compounds  and  measuring  their  quantities  constitute  the  several 
branches  of  analysis.  The  subdivisions  of  chemistry  of  this  kind  are 
numerous,  e.g.,  industrial  chemistry,  bio-chemistry,  food  chemistry, 
radio-active  chemistry,  and  so  forth. 

The  ideal  in  view  in  thus  classifying  the  content  of  a  science  is  to 
convert  it  into  an  organized  body  of  knowledge.  The  various  ways 
used  to  organize  the  facts  of  a  science  will  be  presented  in  detail  as 


CHEMICAL  PHENOMENA  3 

opportunity  offers.     These  ways  constitute  what  is  called  the  scien- 
tific method. 

It  is  only  by  following  intelligently  the  way  in  which  the  science 
is  manufactured,  step  by  step,  out  of  the  raw  material  furnished  by 
observation  and  experiment,  that  the  student  can  gain  a  sound  foun- 
dation for  more  advanced  work  in  the  same  science,  or  a  mental  train- 
ing broad  enough  in  its  tendency  to  add  measurably  to  his  efficiency 
in  every  other  task.  The  chief  object  of  useful  thought,  no  matter 
whether  the  problem  is  one  of  language,  history,  business,  or  life,  is 
to  organize  isolated  facts  into  knowledge,  and  the  means  of  success- 
fully accomplishing  this  is  the  use  of  the  scientific  method. 

What  Chemistry  Deals  With.  —  Chemistry  is  the  science 
which  deals  with  all  forms  of  matter.  It  considers  the  natural  kinds, 
such  as  rocks  and  minerals,  as  well  as  materials  like  fat  and  flour 
obtained  from  animals  or  plants.  It  deals  also  with  artificial  prod- 
ucts like  paints  or  explosives.  When  we  wish  information  about 
any  specimen  or  kind  of  matter,  we  consult  a  chemist.  Now  chem- 
ists have  worked  out  a  point  of  view  which  enables  them  to  attack 
any  problem  'connected  with  matter  in  a  systematic  manner  and  to 
state  the  results  in  a  clear  and  simple  way.  To  learn  chemistry, 
we  must  first  strive  to  acquire  this  point  of  view  and  to  learn  the 
technical  language  the  chemist  uses  in  stating  and  discussing  his 
results. 

Properties.  —  Suppose  that  a  piece  of  rusty  iron  is  submitted 
to  the  chemist.  He  at  once  examines  the  rust  and  notes  that  it  is 
reddish-brown  in  color  and  earthy  in  appearance.  He  separates 
some  of  it  from  the  iron  and  finds  it  to  be  brittle,  that  is,  easily  broken 
and  capable  of  being  pulverized  in  a  mortar.  He  finds  that  its  den- 
sity is  about  4.5,  that  is  to  say,  1  c.c.  (Appendix  I)  of  it  weighs  about 
4.5  g  (1  c.c.  of  water  at  4°  weighs  1  g.).  On  heating  some  of  it  in  a 
flame,  he  finds  that  it  does  not  melt,  and  must,  therefore,  have  a 
very  high  melting-point.  These  qualities  he  calls  properties,  and 
more  especially  physical  properties.  Since  all  specimens  of  iron- 
rust  show  exactly  the  same  properties,  he  often  calls  them  specific 
physical  properties,  because  they  are  properties  shown  by  all  speci- 
mens of  a  particular  species  of  matter. 

After  removing  any  rust  by  filing  or  scraping,  the  chemist  ex- 
amines the  iron,  and  finds  a  fresh,  clean  surface  to  be  almost  white 
and  metallic  in  appearance.  The  metal  is  tenacious,  so  that  it  can 


4  INORGANIC  CHEMISTRY 

be  bent  but  not  easily  broken.  He  finds  that  its  density  is  about 
7.5,  and  that  the  metal  is  incapable  of  being  melted  in  an  ordinary 
flame.  In  addition,  he  finds  it  to  be  strongly  attracted  by  a  magnet, 
while  rust  is  not  attracted. 

The  chemist,  then,  studies  what  he  calls  the  specific  physical 
properties  of  each  material,  in  order  that  he  may  be  able  to  recognize 
various  materials. 

Substances.  —  All  specimens  of  iron  show  one  set  of  proper- 
ties and  all  specimens  of  iron-rust  show  a  different  set,  peculiar  to 
rust.  The  chemist  calls  any  definite  variety  of  matter,  all  speci- 
mens of  which  show  the  same  properties,  a  substance.  Iron  is 
one  substance  and  rust  another.  A  substance  is  recognized  by  its 
properties. 

That  all  bodies  of  a  like  kind  have  identical  properties  is  the 
most  fundamental  fact  in  chemistry.  This  fact  is  called  the  Law 
of  specific  physical  properties,  and  is  stated  thus :  The  specific  physical 
properties  of  a  substance  are  constant  in  all  specimens. 

The  point  of  view  of  the  chemist  consists,  therefore,  in  describing 
any  material  by  ascertaining  whether  it  is  made  of  one,  or  of  more 
than  one  substance.  He  describes  it  by  naming  the  substances  which, 
by  a  study  of  their  properties,  he  has  found  in  it. 

The  foregoing  definition  of  a  substance  is  very  incomplete.  Thus,  a  three 
per  cent  solution  of  common  salt  in  water  always  has  the  same  properties,  yet  it  is 
not  one  substance,  but  a  solution  containing  several  substances.  This  definition 
is  sufficient  for  the  present  purpose,  however. 

Two  Illustrations  of  the  Study  and  Description  of  Mate- 
rials. —  If  a  piece  of  granite  is  examined  by  a  chemist,  he  observes 
at  once  that  it  is  spotted  in  appearance,  and  made  up  of  several 
crystalline  materials  of  differing  nature.  He  therefore  breaks  it 
up  and  studies  the  properties  of  the  fragments.  Some  of  the  frag- 
ments of  granite  are  dark  and  with  a  penknife  can  easily  be  split  into 
transparent  sheets,  thinner  than  paper.  These  particular  frag- 
ments are  in  all  respects  like  mica  (Fig.  1).  This  substance  is  a 
mineral  which,  in  certain  neighborhoods,  occurs  in  large  masses,  and 
sheets  of  it  (" isinglass")  are  used  to  close  the  windows  of  stoves. 
Others  of  the  fragments  are  clear  like  glass,  and  are  very  hard  (see 
Appendix  II),  and  have  all  the  properties  of  quartz  or  rock  crystal 
(Fig.  2),  which  is  another  substance  well  known  to  the  chemist. 
The  remaining  fragments  are  less  clear  than  quartz,  and  are  not  so 


CHEMICAL  PHENOMENA  5 

hard.  They  can  be  split  into  layers,  but  not  nearly  so  easily  as  can 
mica.  They  form  oblong  crystals,  differing  in  this  also  from  quartz, 
which  shows  hexagonal  crystals.*  This  substance  is  felspar  (Fig.  3). 
Thus  the  chemist  studies  the  physical  properties  of  the  fragments, 
and  finds  that  there  are  three  different  substances  in  granite.  He 
reports  that  the  components  of  granite  are  mica,  quartz,  and  felspar. 


FIG.  1. 


FIG.  2. 


FIG.  3. 


When  flour  is  examined  by  the  chemist,  it  appears  to  the  eye  to 
be  all  alike.  Under  the  microscope,  even,  all  he  can  learn  is  that  it 
consists  largely  of  grains,  which  have  the  characteristic  appearance 
(first  property)  of  grains  of  starch  (Fig.  4).  He  places  some  flour 
on  a  square  piece  of  cheese-cloth  and  encloses  it  by  tying  with  a 
thread  (Fig.  5).  On  kneading  the  little  bag  in  a  vessel  of  water, 


FIG.  4. 


FIG.  5. 


milky  water  is  squeezed  out.  When  the  milky  water  stands,  the 
white  material  settles  to  the  bottom,  the  water  can  be  poured  off, 
and  the  deposit  can  be  dried.  This  white  substance,  when  boiled 
with  water,  gives  an  almost  clear  liquid  which  jellies  on  cooling. 
This  is  another  property  of  starch.  A  little  tincture  of  iodine  (solu- 

*  Crystals  (see  also  Index)  are  natural  forms  of  geometrical  outline,  which 
solid  substances  assume.  Usually  each  substance  has  a  more  or  less  distinct 
form  of  its  own,  the  particular  angles  at  which  the  faces  meet  being  peculiar  to 
the  substance.  Its  individual  crystalline  form  is  therefore  a  specific  physical 
property  of  each  substance. 


6  INORGANIC  CHEMISTRY 

tion  of  iodine  in  alcohol),  dropped  on  a  part  of  the  starch,  causes 
the  latter  to  turn  blue.  This  is  a  very  characteristic  property  of 
(and  therefore  test  for)  starch.  When  the  bag  of  flour  is  kneaded 
persistently  in  water  which  is  frequently  changed,  the  material 
finally  ceases  to  render  the  water  milky.  The  starch  has  all  been 
washed  out.  When  the  bag  is  now  opened,  a  sticky  material  is 
found  in  it.  This  is  called  gluten.  The  chemist  therefore  finds  that 
the  flour  contains  starch  and  gluten.  He  learns  this  by  separating 
the  components. 

Law  of  Component  Substances.  —  Every  material  can  be  de- 
scribed as  being  composed  of  one  substance,  or  as  being  a  mixture  of 
two  or  more  component  substances,  each  of  which  has  a  definite  set 
of  specific  physical  properties.  This  is  the  second  fundamental  law 
of  chemistry.  This  conception  was  first  clearly  stated  by  Lomo- 
nossov  (1742),  a  Russian  author,  statesman,  and  chemist  (1711- 
1765). 

Mixtures  and  Impurities.  —  A  material  containing  more  than 
one  component  substance  is  called  a  mixture.  The  characteristic 
of  a  mixture  is  that  each  of  the  component  substances,  although 
mixed  with  the  others,  possesses  exactly  the  same  properties  as  if 
it  were  present  alone.  No  one  of  the  components  affects  any  other 
component  or  alters  any  of  its  properties.  Granite  and  flour  are 
typical  mixtures. 

When  a  specimen  is  composed  mainly  of  one  substance,  and 
contains  only  minute  amounts  of  one  or  more  other  substances,  it 
is  frequently  spoken  of  as  a  specimen  of  the  main  substance  contain- 
ing certain  specified  substances  as  impurities.  To  be  called  an 
impurity,  the  foreign  matter  need  not  be  dirty  or  offensive.  Thus, 
common  salt  usually  contains  a  little  magnesium  chloride,  a  white 
crystalline  solid,  as  an  impurity,  and  it  is  this  impurity  which  becomes 
damp  in  wet  weather.  Again,  compounds  of  lime  and  of  mag- 
nesium are  common  impurities  in  drinking  water. 

Chemically  pure  (C.P.)  means  that  the  quantities  of  the  im- 
purities which  the  material  is  most  apt  to  contain  have  been  reduced 
below  the  amount  which  would  interfere  with  the  most  exact  chemical 
work  for  which  the  substance  is  commonly  employed.  Absolutely 
pure  bodies  are  unknown. 

By  convention  we  continually  speak  of  "pure"  hydrochloric  acid, 
or  of  "pure"  sulphuric  acid,  although  there  may  be  more  than  60  per 


CHEMICAL  PHENOMENA  7 

cent  of  water  present  in  the  former,  and  7  per  cent  in  the  latter.  By 
this  we  mean  to  distinguish  the  former  from  "commercial"  hydro- 
chloric acid,  for  example,  which  contains,  in  addition  to  the  water, 
impurities  like  sulphuric  acid  and  a  coloring  matter.  The  water 
is  in  fact  disregarded,  since  it  is  assumed  to  be  present  in  all  cases. 

Components.  —  The  ingredients  of  a  mixture  are  called  the  com- 
ponents (Lat.,  componere,  to  put  together),  because  they  are  simply 
placed  together,  without  change,  and  can  be  separated  without 
change. 

Bodies  or  Specimens.  It  will  be  seen  that  substance  is  a  gen- 
eral term,  like  the  word  "dog,"  covering  the  whole  species.  The 
substance  iron  includes  all  the  iron  in  the  Universe,  and  all  that  was 
made  in  the  past  or  may  be  made  in  the  future.  When  we  refer 
to  a  particular  piece  of  iron,  we  call  it  a  body  or  a  specimen. 

A  body  is  any  particular  specimen  of  matter,  such  as  a  piece  of 
sulphur,  a  portion  of  water,  a  piece  of  ferrous  sulphide,  a  fragment 
of  granite,  or  some  nitrate  of  silver  solution.  There  are  thus  as 
many  bodies  as  there  are  discrete  portions  of  matter.  A  body  may 
be  heterogeneous,  or  made  up  of  visibly  unlike  parts,  as  granite  and  a 
mixture  of  iron  powder  and  sulphur  are;  or  it 
may  be  homogeneous,  or  alike  in  all  parts,  as 
are  pieces  of  sulphur  and  of  ferrous  sulphide 
and  portions  of  water  and  of  nitrate  of  silver 
solution. 

The  Rusting  of  Metals.  —  If  we  return 
once  more  to  the  subject  of  rusty  iron,  we 
find  another  point  which  interests  the  chemist. 
If  the  iron  is  kept  moist  —  for  example,  by 
lying  in  the  grass  or  partly  immersed  in  water 
—  the  layer  of  rust  gradually  becomes  thicker, 
and  the  core  of  iron  becomes  thinner  until  it 
finally  disappears.  The  rust  seems  to  be 
formed  from  the  iron,  in  presence  of  air  and  Fla*  6* 

moisture.  The  iron,  particle  by  particle,  loses  the  properties  of 
iron  and  simultaneously  acquires  those  of  rust.  Now,  the  chemist 
is  concerned,  not  only  with  recognizing  substances,  but  also  with 
the  ways  in  which  substances  change  and  new  substances  are 
produced. 


8  INORGANIC  CHEMISTRY 

Several  other  metals  rust,  as  does  iron,  but  the  change  is  slower. 
Thus,  lead  rusts  (tarnishes)  slowly,  and  zinc  still  more  slowly.  The 
change  can  be  hastened  by  heating.  If  some  lead  be  melted  in  a 
porcelain  crucible  (Fig.  6)  and  be  stirred  with  an  iron  wire,  a  dirty 
yellow  powder  collects  on  the  surface.  Gradually  more  and  more 
of  the  powder  is  formed  and  less  and  less  of  the  metallic  lead  re- 
mains, until  at  last  all  the  metal  is  gone.  Melted  tin,  when  treated 
in  the  same  way,  gives  a  white  powder. 

Explanation  of  Rusting.  —  The  first  fact  which  seemed  to 
throw  light  on  the  subject  was  discovered  by  a  French  physician, 
Jean  Rey  (1630),  who  found  that  the  rusts  of  tin  and  lead,  made  by 
heating  and  stirring,  were  heavier  than  the  original  pieces  of  metal. 
He  inferred,  correctly,  that  the  additional  material  which  caused 
the  increase  in  weight  came  from  the  air.  He  imagined,  however, 
that  the  rust  was  not  a  new  substance,  but  a  sort  of  froth,  and  there- 
fore a  mixture  of  air  with  the  metal.  Other  investigators,  such  as 
Hooke  (1635-1703)  and  particularly  Mayow  (1645-1679),  in  Eng- 
land, explained  the  increase  in  weight  by  supposing  that  some 
material  from  the  air  had  combined  with  the  metal.  In  other  words, 
iron,  for  example,  was  one  substance  composed  of  iron  only,  and  rust 
was  another  substance,  made  by  union  of  iron  and  a  material  from  the 
air,  and  not  a  mere  mixture. 

It  was  Lomonossov  (1756)  who  first  proved  by  an  experiment 
that  the  extra  material  did  come  from  the  air.  He  placed  some 
tin  in  a  flask,  sealed  up  the  mouth  of  the  vessel,  and  weighed  the 
whole.  The  flask  was  then  heated  and  the  tin  was  converted  into 
the  white  powder.  So  long  as  the  flask  remained  sealed,  no  change 
in  weight  was  found  to  have  occurred.  When  the  mouth  of  the 
flask  was  opened,  however,  some  air  rushed  in,  and  the  total  weight 
was  then  found  to  be  greater.  Evidently  a  portion  of  the  original 
air,  during  the  heating,  had  forsaken  the  gaseous  condition  and 
joined  itself  to  the  tin  to  form  the  powder.  This  left  a  partial 
vacuum  in  the  flask,  and  more  air  entered  when  the  latter  was 
opened.  Eighteen  years  later  the  same  experiment  was  made  by 
Lavoisier,  who  drew  trie  same  conclusion.  The  rusting  of  other 
metals  was  found  to  be  due  to  the  same  cause.  Lavoisier  named 
the  gas,  taken  from  the  air,  oxygen. 

The  conclusion  can  be  confirmed  in  various  ways.  For  example, 
when  the  air  is  pumped  out  of  the  flask  before  it  is  sealed,  the  metal 
can  be  heated  in  the  vacuum  indefinitely  without  rusting. 


CHEMICAL  PHENOMENA 


9 


FIG.  7. 


A  rough  imitation  of  the  rusting  of  iron  may  be  shown  to  be  accompanied  by 
an  increase  in  weight.  Iron  powder  is  suspended  by  means  of  a  magnet  over  one 
pan  of  a  balance,  and  the  equipoise  is  restored  by  placing  small  shot  on  the  other 
pan.  When  the  iron  is  heated,  union  with  oxygen  begins  and,  after  a  time,  the 
pointer  inclines  markedly  to  one  side.  The  product  here  is  magnetic  oxide  of 
iron  Fe3O4,  however,  and  not  rust  Fe2O3,xH2O. 

Experiment  to  Show  the  Nature  of  Rusting.  —  That  a  part 
of  the  air  is  consumed  when  iron  rusts  is  easily  proved.  We  moisten 
the  interior  of  a  test-tube  and  sprinkle  some  powdered  iron  so  that 
it  covers  and  adheres  to  the  whole  interior 
surface.  We  then  set  the  tube  mouth  down- 
wards in  a  dish  of  water  (Fig.  7).  At  first, 
the  pressure  of  the  water  compresses  the  air  in 
the  tube  very  slightly,  and  the  water  ascends 
above  the  mouth  to  the  extent  of  a  small  frac- 
tion of  an  inch  only.  As  the  moist  iron  slowly 
rusts,  however,  the  oxygen  is  gradually  re- 
moved, and  the  pressure  of  the  atmosphere 
outside  slowly  pushes  the  water  further  up  the 
tube.  After  an  hour  or  more,  the  water  has 
ascended  about  one-fifth  of  the  total  distance  towards  the  top  of  the 
tube.  Evidently  part  of  the  air  has  changed  from  the  gaseous  con- 
dition, and  the  water  has  been  forced  up  to  take  its  place.  Inspec- 
tion now  shows  some  reddish  particles,  where  rusting  has  taken  place. 
The  rust,  then,  contains  part  of  the  iron  and  all  the  oxygen  that 
the  tube  contained. 

Of  course,  much  of  the  iron  powder  is  still  grey,  and  has  not 
rusted.  The  air  in  the  tube  did  not  contain  oxygen  enough  to  com- 
bine with  all  the  iron.  The  iron  that  remains  is  as  little  able  to 
rust  in  the  remaining  gas  as  in  a  vacuum. 

Incidentally  we  learn  from  this  experiment  that  atmospheric  air 
contains  about  one-fifth  (20  per  cent)  oxygen  by  volume.  The  re- 
maining four-fifths  is  almost  all  nitrogen  (79  per  cent),  a  substance 
which  combines  with  very  few  materials,  while  the  balance  (1  per 
cent)  is  made  up  of  gases  which  do  not  enter  into  combination  with 
any  known  substance.  If  lead,  tin,  or  zinc  had  been  heated  in  an 
enclosed  volume  of  air,  they  likewise  would  have  taken  out  the 
20  per  cent  of  oxygen  and  would  have  left  the  other  gases. 

An  Older  View:  Phlogiston.  —  Simple  and  satisfactory  as  this 
explanation  appears  to  us,  it  must  be  said  that  it  gained  little  sym- 


10  INORGANIC  CHEMISTRY 

pathy  from  the  contemporaries  of  Rey,  Hooke,  and  Mayow.  The 
other  investigators  had  been  prejudiced  by  a  remarkable  hypothesis 
which  was  supposed  to  explain  both  rusting  and  combustion.  Start- 
ing with  a  suggestion  of  Plato's,  that  during  combustion  some  material 
escaped  from  the  burning  body,  and  that  the  flames  and  heat  repre- 
sented the  vigor  with  which  this  substance  rushed  out,  Stahl  and 
Becher  invented  the  idea  that  a  substance,  which  they  called  "  phlo- 
giston," was  contained  in  all  materials  capable  of  rusting  or  burning. 
Its  escape  accounted  for  the  phenomena  of  combustion,  and  its  ab- 
sence for  the  alteration  in  properties  of  the  residual  substance.  They 
were  perfectly  aware  that  the  material  was  heavier  after  rusting  than 
before,  but  refused  to  sacrifice  their  hypothesis  to  a  mere  fact  like 
this.  So  they  ingeniously  appended  the  suggestion  that  phlogiston 
was  a  substance  which  not  only  was  not  subject  to  gravitation,  but 
possessed  the  opposite  property  of  levity !  Thus,  its  escape  rendered 
the  material  from  which  it  issued  heavier  than  before!  Instead  of 
demanding  the  preparation  and  examination  of  phlogiston  itself,  and 
the  demonstration  that  it  weighed  less  than  nothing,  the  generality  of 
chemists  of  that  age  accepted  the  idea  without  proof,  It  is  not  sur- 
prising, therefore,  that  many  of  their  attempts  to  explain  chemical 
experiments  on  the  basis  of  an  assumption  which  was  the  precise 
opposite  of  the  truth  should  have  resulted  in  hopeless  confusion. 
The  fact  that  foreign  matter  was  actually  gained  by  a  body  during  the 
process  of  rusting  was  not  generally  accepted  until  it  was  demon- 
strated anew  by  Lavoisier  in  1774  (p.  8).  The  whole  development 
of  chemistry  was  stunted  by  the  general  belief  in  the  conception  of 
phlogiston  and,  during  the  one  hundred  and  fifty  years  which  passed 
between  the  work  of  Jean  Rey  and  his  contemporaries,  and  that  of 
Lavoisier,  relatively  little  progress  was  made. 

The  introduction  of  the  balance  into  the  chemical  laboratory,  and  the  first 
use  of  measurements  of  weight  as  a  means  of  exploring  and  explaining  chemical 
changes,  is  frequently  ascribed  to  Lavoisier.  As  a  matter  of  fact,  it  is  difficult  to 
state  when  measurement  of  weight  first  became  the  chief  ally  of  the  chemist  in  his 
work.  Important  and  conclusive  results  were  obtained  by  its  means,  however, 
before  the  time  of  Lavoisier,  for  example,  by  Jean  Rey,  Boyle,  Lomonossov,  and 
Black  (1728-1799). 

Explanation  in  Science.  —  One  section  (p.  8)  was  entitled 
Explanation  of  rusting.  If  that  paragraph  be  now  re-read,  it  will 
be  found  that,  in  the  ordinary  (as  distinct  from  the  scientific)  sense 
of  the  word,  no  explanation  was  given!  When  we  ask  a  man  to  "ex- 


CHEMICAL  PHENOMENA  11 

plain"  some  feature  in  his  conduct,  we  recognize  that  he  might 
have  chosen  to  act  otherwise,  and  we  wish  to  know  why  he  acted 
precisely  as  he  did.  Nature,  however,  has  no  free  will,  and  cannot 
tell  why  she  presents  certain  phenomena,  and  not  others. 

On  examining  the  explanation,  we  find  that  it  simply  shows  that, 
when  iron  rusts,  it  combines  with  oxygen  from  the  air.  This  is  an 
additional  fact.  It  shows  how  iron  rusts,  namely  by  taking  up 
oxygen,  but  not  why  it  is  able  to  unite  with  oxygen.  We  simply 
do  not  know  why  iron  can  combine  with  oxygen  gas  and  platinum 
can  not. 

Explanations  in  chemistry  are  of  three  kinds.  (1)  We  usually 
try  to  show  that  the  phenomenon  is  not  an  isolated  one.  Thus, 
we  show  that  other  metals  rust.  This  reconciles  us  to  some  extent 
to  the  fact  that  iron  rusts,  and  we  feel  some  mental  satisfaction. 
This  is  the  method  of  showing  that  the  fact  to  be  explained  is  a 
member  of  a  large  class  of  similar  facts.  (2)  Next  we  try  to  get 
more  information  about  the  fact  to  be  explained.  Thus,  when,  to 
the  acquaintance  with  the  outward  manifestations  of  rusting,  we 
add  the  further  information  that  there  is  an  increase  in  weight,  and 
that  this  is  due  to  union  of  oxygen  from  the  air  with  the  iron,  we  feel 
increased  satisfaction,  and  say  that  the  fact  has  been  ' 'explained." 
(3)  If  we  are  still  dissatisfied,  and  can  discover  no  further  useful 
facts,  we  imagine  a  state  of  affairs  which,  if  true,  would  classify  the 
fact  or  add  to  what  we  know  about  it.  This  step  we  call  explain- 
ing by  means  of  an  hypothesis.  We  then  devote  our  attention  to 
trying  to  verify  the  hypothesis.  The  making  of  attempts  to  explain 
facts  is  part  of  the  Scientific  Method  (p.  3). 

The  Law  of  Chemical  Change.  —  The  three  examples  of  rust- 
ing show  that  specimens  of  matter  can  lose  their  original  properties 
and  acquire  new  ones.  Since  a  substance  is  "a  species  of  matter, 
with  a  constant  set  of  properties,"  we  are  compelled  to  decide  that, 
when  a  material  changes  its  properties,  it  has,  in  doing  so,  become  a 
new  substance.  This  consideration  calls  to  our  attention  the  third 
of  the  fundamental  laws  of  chemistry,  namely,  that  the  material 
forming  one  or  more  substances  (such  as  oxygen  and  iron),  without 
ceasing  to  exist,  may  be  changed  into  one  or  more  entirely  different 
substances.  Such  a  change  is  called  a  chemical  change,  or  action, 
or  interaction,  or  reaction. 

The  commoner  kinds  of  chemical  actions  can  be  divided,  for  con- 
venience, into  five  varieties.  We  can  now  define  the  first  of  these. 


12  INORGANIC  CHEMISTRY 

First    Variety    of    Chemical    Change:     Combination.  —  In 

each  case  of  rusting,  two  substances  (a  gas  and  a  metal)  come  to- 
gether to  form  a  third  substance  (an  earthy  powder).  Apparently 
two  substances  may  come  together  in  two  different  ways.  They 
may  form  a  mixture,  in  which  both  substances  are  present,  and 
retain  their  properties,  or  they  may  come  together  to  form  a  single 
substance  with  different  properties.  t  When  two  (or  more)  sub- 
stances unite  to  form  one  substance,  the  change  is  called  chemical 
combination  or  union.  The  product  is  called  a  compound  substance. 
We  are  very  careful  never  to  speak  of  a  compound  substance  as  a 
mixture.  Rust  is  not  a  mixture  of  iron  and  oxygen;  it  shows  none 
of  the  properties  of  either.  Nor  do  we  call  a  mixture  (like  granite)  a 
compound,  or  the  operation  of  mixing,  combination  or  union.  These 
are  technical  words  in  chemistry  and,  to  avoid  confusion,  may  be 
used  only  with  due  regard  to  their  technical  meanings. 

Constituents.  —  As  we  have  seen,  we  speak  of  the  substances 
in  a  mixture  as  the  components.  When  we  wish  to  refer  to  the  forms 
of  matter  which  are  chemically  united  in  a  compound,  we  call  them 
the  constituents  (Lat.,  con  and  statuere,  to  stand  together)  of  the 
compound  substance.  Thus,  iron  and  oxygen  are  the  constituents 
of  rust. 

The  chemist  separates  (p.  6)  the  components  of  a  mixture,  for 
that  is  all  that  is  necessary.  He  liberates  the  constituents  of  a  com- 
pound, however,  because  they  are  bound  together  in  chemical  com- 
bination. 

The  names  given  to  compounds  are  usually  devised  so  as  to 
indicate  the  nature  of ,  the  constituents.  Thus,  iron-rust  is  oxide 
of  iron  (or  ferric  oxide,  from  Lat.  ferrum,  iron).  The  yellowish 
powder  from  lead  is  lead  oxide  or  oxide  of  lead,  and  the  white  powder 
from  tin  is  oxide  of  tin. 

A  Condensed  Form  of  Statement.  —  We  may  represent  a 
chemical  combination,  or  indeed  any  kind  of  chemical  change,  in  a 
condensed  form,  thus: 

Iron  -f  Oxygen  — >  Oxide  of  iron  (ferric  oxide). 

Each  name  stands  for  a  substance.  Two  substances  in  contact  with 
one  another  (mixed),  but  not  united  chemically,  are  connected  by 
the  +  sign.  The  arrow  shows  where  the  chemical  change  comes  in, 
and  the  direction  of  the  change.  We  read  the  statement  thus: 


CHEMICAL  PHENOMENA  13 

Iron  and  oxygen  brought  together  under  suitable  conditions  undergo 
chemical  change  into  oxide  of  iron,  called  also  ferric  oxide.  Similarly 
we  may  write: 

Lead  +  Oxygen  — » Oxide  of  lead. 

Tin    +  Oxygen  — »  Oxide  of  tin. 

The  Increase  in  Weight  in  Rusting.  —  As  we  have  seen,  the 
process  of  rusting  is  accompanied  by  a  slow  increase  in  th3  weight 
of  the  solid,  due  to  the  gradual  addition  of  oxygen  to  the  metal. 
Now,  this  increase  in  weight  ceases  of  its  own  accord,  when  a  certain 
maximum  has  been  reached.  This  occurs  when  the  last  particles  of 
the  metal  have  disappeared.  Thus,  the  lead  gains  in  weight  until 
every  100  parts  of  the  metal  have  gained  7.72  parts  of  oxygen,  and 
the  tin  until  every  100  parts  have  gained  26.9  parts  of  oxygen. 
When  these  increases  have  occurred,  the  metal  is  found  to  have  been 
all  used  up,  and  prolonged  heating  and  stirring  cause  no  further 
absorption  of  oxygen  and  no  further  change  in  weight.  This  fact, 
that  each  substance  limits  itself  of  its  own  accord  to  combining  with  a 
fixed  proportion  of  the  other  substance,  in  forming  a  given  compound, 
is  one  of  the  most  striking  facts  about  chemical  combination.  In 
mixtures,  any  proportions  chosen  by  the  experimenter  may  be  used. 
In  chemical  union,  the  experimenter  has  no  choice;  the  proportions 
are  determined  by  the  substances  themselves.  Thus,  100  parts  of 
iron,  when  turning  into  ordinary,  red  rust,  take  up  43  parts  of  oxygen, 
no  more  and  no  less. 

This  fact  enables  us  to  make  our  condensed  statements  more 
specific  and  complete  by  including  in  them  the  proportions  by  weight 
used  in  the  chemical  change: 

Iron    (100)  +  Oxygen  (43)  ->  Ferric  oxide  (143). 
Lead  (100)  +  Oxygen  (7.72)  ->  Oxide  of  lead  (107.72). 

The  following  numbers,  which  represent  the  same  proportions 
by  weight,  are  the  ones  commonly  used  by  chemists: 

Iron  (111.68)  +  Oxygen  (48)  ->  Ferric  oxide  (159.68). 

Summary.  —  Thus  far,  we  have  learned  that  chemistry  deals 
with  substances  and  their  physical  properties,  and  with  the  changes 
which  substances  undergo.  We  have  discussed  and  defined  a  number 
of  important  words  expressing  fundamental  chemical  ideas.  Finally, 
we  have  touched  upon  the  weights  of  the  materials  used  in  chemical 


14  INORGANIC  CHEMISTRY 

change,  a  subject  of  great  importance  which  will  be  developed 
further  in  a  later  chapter. 

We  must  now  take  up  four  new  examples  of  chemical  change. 
They  will  aid  us  in  introducing  one  or  two  additional  conceptions 
and  laws  that  are  continually  used  by  the  chemist,  and  without 
which  we  cannot  begin  the  systematic  study  of  the  science. 

Another  Case  of  Combination:  Iron  and  Sulphur.  —  Since 
oxygen  is  an  invisible  gas,  there  is  a  slight  difficulty  in  realizing  that 
rusting  consists  in  the  union  of  two  substances  —  this  gas  and  a 
metal.  The  present  example  is  less  interesting  historically,  but  it 
is  simpler  because  both  substances  are  visible  and  are  easily  handled. 
The  case  of  iron  and  sulphur  will  enable  us  to  illustrate  the  same 
point  of  view  and  to  practice  the  application  of  the  same  technical 
words.  It  will  also  introduce  us  to  two  manip- 
ulations—  filtration  and  evaporation  —  which 
are  frequently  used  by  the  chemist. 

We  begin  by  observing  the  physical  proper- 
ties of  the  two  substances.     Those  of  iron  have 
already  been  noted  (pp.  3-4)  .*    Sulphur  is  a  pale- 
yellow  substance  of  low  density  (p.  3),  namely, 
FIQ  g  2.     It  is  easily  melted  (m.-p.  112.8°).     It  does 

not  dissolve  in  water  —  that  is,  it  does  not  mix 

completely  with  and  disappear  in  water,  as  sugar  does  on  stirring. 
It  does  dissolve  readily  in  carbon  disulphide,  however  (41  parts:  100 
at  18°).  It  crystallizes  in  rhombic  forms  (Fig.  8).  It  is  not  at- 
tracted by  a  magnet. 

Study  of  the  Mixture,  before  Combination.  —  Now,  if  some 
iron  filings  and  pulverized  sulphur  are  stirred  together  in  a  mortar, 
the  result  is  a  mixture.  True,  the  color  is  not  that  of  either  sub- 

*  References  to  previous  pages  are  used  in  order  to  save  needless  repetition 
in  writing.  The  beginner,  however,  requires  endless  repetition  in  his  reading 
and  must  form  the  habit  of  examining,  in  conjunction  with  the  current  text,  the 
parts  referred  to.  The  passages  cited  are,  by  the  reference,  made  part  of  the 
current  text,  which  will  usually  not  be  clear  without  them.  The  same  remark 
applies  to  topics  referred  to  by  name.  Such  topics,  if  treated  in  later  pages, 
are  distinguished  by  the  letters  q.v.  (quod  vide  —  which  see),  and  must  be  sought 
in  the  index. 

All  terms,  and  especially  those  borrowed  from  physics,  if  not  perfectly 
familiar,  must  be  looked  up  in  a  work  on  physics  or  in  a  dictionary. 


CHEMICAL  PHENOMENA 


15 


FIG.  9. 


stance,  but  with  a  lens  particles  of  both  substances  can  be  seen. 
Passing  a  magnet  over  the  mixture  will  easily  remove  a  part  of  the 
iron,  and  with  the  help  of  a  lens  and  a  needle  the 
mixture  can  be  picked  apart,  particle  by  particle, 
completely.  We  can  separate  the  components  of  the 
mixture  more  expeditiously,  however,  by  using  ma- 
nipulations based  upon  certain  suitable  properties. 
Thus,  sulphur  dissolves  in  carbon  disulphide  while 
iron  does  not.  If,  therefore,  a  part  of  the  mixture 
is  placed  in  a  dry  test-tube  along  with  some  carbon 
disulphide  (Fig.  9)  and  is  shaken,  the  liquid  dissolves 
the  sulphur  and  leaves  the  iron.  To  complete  the 
separation,  the  iron  must  be  removed  from  the  liquid 
by  nitration,  and  the  sulphur  recovered  by  evaporation  of  the  car- 
bon disulphide. 

Filtration.  —  Iron,  or  any  solid,  when  mixed  with  a  liquid  or 
solution  (like  the  solution  of  sulphur  in  carbon  disulphide)  is  said  to 
be  suspended  in  the  liquid.  If  the  solid  is  one  that  settles  rapidly, 

the  liquid  may  be  separated  from  the  solid 
in  a  rough  way  by  pouring  off  as  much  of 
the  clear  liquid  as  possible.  This  is  called 
decantation. 

A  complete  separation  is  best  made  by 
pouring  the  mixture  on  to  a  cone  of  filter 
paper  supported  in  a  glass  funnel  (Fig.  10). 
The  liquid,  together  with  anything  that 
may  be  dissolved  in  it,  runs  through  the 
pores  of  the  paper  and  down  the  hollow 
stem  of  the  funnel.  The  liquid  is  then 
called  the  nitrate.  The  particles  of  the  sus- 
pended solid  are  too  large  to  pass  through 
the  pores,  and  so  collect  on  the  surface  of 
the  filter  paper.  This  operation,  like  everything  the  chemist  does, 
takes  advantage  of  the  physical  properties  of  the  various  materials. 
The  material  remaining  on  the  paper  (the  residue),  when  dry,  is 
wholly  attracted  by  a  magnet  and  shows  all  the  other  properties  of 
iron. 

Evaporation.  —  To  recover  the  sulphur,  the  solution  in  carbon 
disulphide  —  the  filtrate  —  is  poured  into  a  porcelain  evaporating 
dish  (Inflammable!  Keep  flames  away).  When  the  vessel  is  set 


FIG.  10. 


16  INORGANIC  CHEMISTRY 

aside,  the  liquid  gradually  passes  off  in  vapor  (e-vapor-ates).  The 
sulphur,  however,  gives  practically  no  vapor  at  room  temperature 

and  remains  as  a  residue  in  the  form  of 
crystals  of  rhombic  outline  in  the  bottom 
of  the  dish  (Fig.  11).  Here,  again,  physical 
properties  have  been  utilized. 

Since  the  physical  properties  of  two  sub- 
stances are  not  changed  by  mixing,  we  have 

thus  used  the  properties  of  the  iron  and  sulphur  so  as  to  separate 
them  once  more.  The  iron  is  on  the  paper;  the  sulphur  is  in  the 
dish. 

Combination  of  Iron  and  Sulphur.  —  But  iron  and  sulphur 
are  capable  of  combining.  If  we  alter  the  conditions  by  raising  the 
temperature  of  some  of  the  dry  mixture,  as  we  did  in  causing  lead  to 
rust  rapidly,  chemical  union  sets  in.  When  we  place  some  of  the 
original  mixture  of  iron  and  sulphur  into  a  clean  test-tube  and  warm 
it,  we  soon  notice  a  rather  violent  development  of  heat  taking  place, 
the  contents  begin  to  glow,  and  what  appears  to  be  a  form  of  com- 
bustion spreads  through  the  mass.  The  heating  employed  at  the 
start  falls  far  short  of  accounting  for  the  much  greater  heat  pro- 
duced. When  these  phenomena  have  ceased,  and  the  test-tube  has 
been  allowed  to  cool,  we  find  that  it  now  contains  a  somewhat  porous- 
looking,  black  solid.  This  material  is  brittle;  it  is  not  magnetic; 
it  does  not  dissolve  in  carbon  disulphide;  and  close  examination, 
even  under  a  microscope,  does  not  reveal  the  presence  of  different 
kinds  of  matter.  This  substance  is  known  to  chemists  as  ferrous 
sulphide  and,  as  we  see,  its  properties  are  entirely  different  from 
those  of  the  constituents. 

In  this  connection  we  must  not  omit  to  notice  that,  as  in  rusting,  a 
certain  fixed  proportion  will  be  used  in  forming  the  compound.  We 
find  that,  for  7  parts  of  iron,  almost  exactly  4  parts  by  weight  of 
sulphur  are  required.  If  more  iron  is  put  into  the  original  mixture, 
then  some  unused  iron  will  be  found  in  the  mass  after  the  action. 
If  too  much  sulphur  is  employed,  some  may  be  driven  off  as  vapor 
by  the  heat  and  all  that  remains,  beyond  the  correct  proportion, 
can  be  dissolved  out  of  the  ferrous  sulphide  with  carbon  disulphide. 
The  sulphur  which  has  combined  with  the  iron,  however,  is  no  longer 
present  as  sulphur  —  it  has  no  longer  the  properties  of  sulphur,  and 
therefore  cannot  be  dissolved  out: 

Iron  (55.84)  +  Sulphur  (32.06)  -» Ferrous  sulphide  (87.9). 


CHEMICAL  PHENOMENA  17 

Another  Illustration:  Mercuric  Oxide.  —  It  has  long  been 
known  that  air  contains  an  active  and  an  inactive  gas.  The  Chinese 
called  them  yin  and  yang,  respectively.  Mayow  (1643-1679)  showed 
that  the  active  gas  caused  rusting,  that  it  was  absorbed  by  paint 
(really  by  the  linseed  oil)  in  "drying,"  that  it  supported  combustion 
of  wood  and  sulphur,  and  that  it  is  necessary  to  life,  being  absorbed 
by  the  blood  from  the  air  entering  the  lungs.  It  was  not  until  later, 
however,  that  a  pure  specimen  of  this  gas  was  obtained,  and  was 
recognized  to  be  a  special  kind  of  gas  different  from 
ordinary  air.  The  gas  (later  to  be  named  oxygen) 
was  made  by  Bayen  from  mercuric  oxide,  a  bright 
red,  rather  heavy  powder.  When  this  substance  is 
heated  (Fig.  12),  we  find  that  a  gas  is  given  off, 
which  is  easily  shown  to  be  different  from  air,  since 
a  glowing  splinter  of  wood  is  instantly  relighted 
on  being  immersed  in  it.  The  gas  is  pure  oxygen. 
We  notice  also  during  the  heating  that  a  sort  of 
mirror  appears  on  the  sides  of  the  tube.  'Apparently 
the  vapor  of  some  metal  is  coming  off  with  the  oxygen  and  con- 
densing on  the  cool  parts  of  the  tube.  As  this  shining  substance 
accumulates  it  takes  the  form  of  globules,  which  may  be  scraped 
together.  It  is,  in  fact,  the  metal  mercury,  or  quicksilver.  If  the 
heating  continues  long  enough,  the  whole  of  the  red  powder  even- 
tually disappears,  and  is  converted  into  these  two  products. 

Mercuric  oxide  (216.6)  — » Mercury  (200.6)  +  Oxygen  (16) 

The  extensive  nature  of  the  change  in  properties  in  this  case  is  evi- 
dent. It  should  also  be  observed  that  continuous  heating  is  required 
to  maintain  this  change  in  operation.  It  differs  markedly  from  the 
iron  and  sulphur  case  m  this  respect.  When  the  flame  is  removed, 
the  evolution  of  oxygen  ceases.  The  significance  of  this  will  appear 
shortly. 

Second  Variety  of  Chemical  Change:    Decomposition.  — 

Bayen's  experiment  introduces  to  us  a  second,  and  very  common 
kind  of  chemical  action.  The  first  variety  was  combination  or 
union  (p.  12).  The  second  is  called  decomposition.  It  consists 
in  starting  with  a  single  substance  (here  mercuric  oxide)  and  split- 
ting it  into  two  (or  more)  substances,  which  differ  in  properties 
from  the  substance  taken  and  from  one  another.  Here,  the  red 
powder  gave  mercury,  a  liquid  metal,  and  oxygen,  a  colorless  gas. 


18  INORGANIC  CHEMISTRY 

Third  Illustration:  Hydrogen  from  an  Acid.  —  If  some  sul- 
phuric acid,  a  liquid,  be  added  to  4-5  times  its  own  volume  of  water, 
the  acid  dissolves,  and  dilute  sulphuric  acid  is  obtained.  When  zinc 
is  added  to  the  cold  mixture,  bubbles  of  a  gas  —  hydrogen  —  form 
on  the  surface  of  the  zinc  and,  when  they  become  large  enough,  they 
rise  to  the  surface  and  break.  The  gas  burns,  when  set  on  fire,  and 
differs,  therefore,  from  air  and  from  oxygen.  The  nature  of  the 
chemical  change  may  be  seen  in  the  following  abbreviated  state- 
ment: 

Zinc  (65.37)  +  Sulphuric  acid  — >  Zinc  sulphate  +  Hydrogen  (2.016) 

Sulphur  (32.06)          Sulphur  (32.06) 
Oxygen  (64.00)  Oxygen  (64.00) 

Hydrogen  (2.016)       Zinc  (65.37) 

If  sufficient  zinc  is  used,  the  action  will  cease  when  all  the  acid,  often 
called  hydrogen  sulphate,  has  been  acted  upon.  The  liquid  can 
thus  be  poured  away  from  the  remaining  zinc  and  can  be  partly 
evaporated.  The  liquid,  when  cold,  then  deposits  crystals  of  zinc 
sulphate.  This  is  a  colorless  substance,  soluble  in  water,  and  there- 
fore invisible  until  so  much  of  the  water  has  been  driven  off  by  boiling 
that  the  quantity  of  water  which  is  left  is  not  sufficient  in  amount  to 
keep  all  the  zinc  sulphate  in  solution. 

This  chemical  change  is  more  complicated  than  are  the  two  pre- 
viously mentioned.  The  hydrogen  leaves  the  sulphuric  acid,  and 
goes  free,  while  the  zinc  takes  its  place,  and  combines  with  the  sulphur 
and  oxygen.  The  zinc  is  said  to  displace  the  hydrogen.  Zinc  is 
able  also  to  displace,  and  liberate,  metallic  copper  from  cupric  sul- 
phate solution,  and  several  other  metals  from  their  compounds. 

Third    Variety    of    Chemical    Change:     Displacement.  - 
When  a  simple   substance  (here,  the  zinc)  and  a  compound  (here, 
sulphuric  acid)  interact,  so  that  another  simple  substance  is  set  free, 
and  the  first  simple  substance  takes  the  place  of  the  second  in  the 
compound,  the  action  is  called  a  displacement. 

Fourth  Illustration:  Salt  and  Silver  Nitrate.  —  The  fourth 
illustration  is  taken  purposely  in  order  to  illustrate  the  variety  of 
ways  in  which  chemical  change  may  be  carried  out.  It  is  the  inter- 
action of  silver  nitrate  and  sodium  chloride  (common  salt).  The 
sut  stances  may  be  recognized  by  the  form  of  the  crystals  of  which 


CHEMICAL  PHENOMENA 


19 


they  consist.  The  latter  is  composed  of  small  cubes  (Fig.  13), 
while  the  former  presents  a  less  familiar  form  geometrically  (Fig.  14). 
Both  substances  are  capable  of  being  dissolved  in  water  and,  for  this 
experiment,  portions  of  each  substance  are  shaken  in  separate 


FIG.  13. 


FIG.  14. 


vessels  with  water,  until  none  of  the  solid  remains.  When  the 
solutions  are  now  poured  together,  we  observe  that  the  clear  liquids 
at  once  become  opaque,  and  that  a  dense  mass  of  white,  solid  material 
appears  suspended  in  the  mixture  (Fig.  15).  Thi?  white  substance 

consists  of  an  extremely  fine  powder  without 

any  observable  crystalline  form.  We  know  at 
once  that  it  must  represent  a  new  substance, 
since  it  would  not  have  appeared  had  it  been 
soluble  in  water  like  the  two  materials  from 
which  it  was  made.  We  continue  adding  the 
one  liquid  gradually  to  the  other  until  no  further 
formation  of  this  solid  takes  place,  and  then 
stop.  By  filtration  (Fig.  10),  we  obtain  the 
insoluble  material  (the  precipitate)  upon  the 
filter  paper,  and  the  clear  liquid  (the  filtrate) 
passes  through  and  is  caught  in  the  vessel 
below. 

We  are  confronted  with  two  possibilities:  either  both  the  original 
materials  have  come  together  to  form  one  white  insoluble  material,  or 
some  other  product  (or  products)  may  be  present  in  addition  to  it.  In 
the  latter  case,  search  must  evidently  be  made  in  the  liquid.  By 
evaporating  the  filtrate  in  a  suitable  vessel  (Fig.  11),  we  find  that  the 
second  assumption  represents  the  fact,  for  a  considerable  quantity  of  a 
white  crystalline  substance  remains.  The  homogeneous  character  of 
this  shows  that  there  was  but  one  product  in  solution,  while  the  same 
property  of  the  precipitate  shows  that  there  are  but  two  products 
altogether. 

The  insoluble  material  is  composed  of  silver  and  chlorine,  and  it 
is  known  as  silver  chloride.  Like  some  other  compounds  of  silver,  it 
darkens  on  exposure  to  light,  turning  first  purple  and  then  brown, 


FIG.  15. 


20  INORGANIC  CHEMISTRY 

and  being  decomposed  by  this  agency  into  its  constituents.  The 
chlorine,  a  gas,  escapes  into  the  air,  and  a  brown  powder  consisting 
of  finely  divided,  metallic  silver  remains.  The  soluble  solid,  ob- 
tained from  the  filtrate,  we  recognize  as  identical  with  a  mineral, 
sodium  nitrate,  which  is  found  in  Chile.  Its  crystals  are  rhombo- 
hedral  (Fig.  69).  They  resemble  cubes  which  have  been  slightly 
distorted  by  pressing  inwards  two  opposite  corners.  This  change 
presents  several  features  which  distinguish  it  from  the  previous  ones : 
it  is  much  more  complex  (see  next  section);  it  takes  place  in  the 
presence  of  water;  it  requires  no  heating  for  its  promotion;  and  the 
change  is  complete  the  instant  the  materials  have  been  mixed,  while 
the  others  required  a  good  deal  of  time  for  their  accomplishment. 

Fourth  Variety  of  Chemical  Change:  Double  Decomposi- 
tion. —  The  last  illustration  may  be  represented  as  follows,  the 
numbers  showing  the  relative  weights  as  before: 

Silver  nitrate  +  Sodium  chloride  — >  Silver  chloride  +  Sodium  nitrate 

Silver  (107.88)  Sodium  (23.00)  Silver  (107.88)  Sodium  (23.00) 

Nitrogen  (14.01)       Chlorine  (35.46)  Chlorine  (35.46)          Nitrogen  (14.01) 

Oxygen  (48.00)  Oxygen  (48.00) 

Here  the  constituents  of  both  of  the  ingredients  become  separated 
(decomposition),  and  the  products  of  this  decomposition  recombined 
cross-wise  (combination).  Thus,  for  example,  the  silver  of  one 
ingredient  combined  with  the  chlorine  of  the  other.  When  two  com- 
pounds interact,  so  that  each  splits  into  two  parts  (called  radicals), 
and  the  parts  exchange  partners  in  recombining,  the  action  is  called 
a  double  decomposition. 

Fifth  Variety  of  Chemical  Change:  Internal  Rearrange- 
ment. —  When  we  encounter  other  chemical  changes,  we  shall  find 
the  extent  of  the  stride  we  have  taken  in  this  critical  analysis  of  a 
few  examples.  It  will  then  appear  that  the  chemical  changes  of 
matter  are  not  nearly  so  various  in  mechanism  as  we  might  have 
anticipated.  In  fact,  there  are  few  changes  which  cannot  be  placed 
in  one  of  the  four  above  described  categories.  Those  that  cannot 
be  so  placed  belong  to  a  fifth  variety  which,  for  the  sake  of  complete- 
ness, may  now  be  mentioned. 

It  occasionally  happens,  especially  in  the  case  of  compounds  of 
carbon  (see  Urea),  that  one  single  kind  of  material  turns  into  another 
single  kind  of  material.  Nothing  is  added  and  nothing  removed,  yet 
the  new  substance  has  different  properties  in  every  respect  from  the 


CHEMICAL  PHENOMENA  21 

old.  Most  of  those  substances  whose  transformation  is  definitely 
assigned  to  this  class  contain  several  constituents,  but  a  rough  notion 
of  this  sort  of  chemical  change  may  be  obtained  by  considering  the 
two  forms  of  phosphorus.  One  of  them  is  pale  yellow  in  appearance, 
easily  melted,  and  very  easily  ignited;  the  other  is  red,  does  not  melt 
on  being  heated,  and  is  difficult  to  ignite.  The  latter  is  made  from 
the  former  by  heating  it  continuously  for  some  hours  in  a  closed  ves- 
sel at  about  300°.  As  no  material  is  taken  up,  the  weight  of  the  sub- 
stance is  unchanged,  and  yet,  when  the  vessel  is  opened,  the  common 
phosphorus  is  found  to  have  turned  into  the  red  variety.  Now,  we 
have  seen  that  when  substances  have  different  properties  they  differ 
also  in  their  constituent  materials.  Carrying  out  this  idea,  the 
hypothesis  (or  suggestion)  has  been  made  that,  since  here  also  the 
properties  change,  there  must  be  some  readjustment  of  the  material, 
even  in  cases  like  this.  Hence,  we  designate  changes  of  this  kind 
internal  rearrangements.  The  composition  of  the  material  is  un- 
altered, so  we  suppose  its  constitution  to  have  become  different. 
If  the  chemists  ever  decide  to  regard  the  change  of  water  into  ice  or 
steam  as  a  chemical  phenomenon,  all  changes  of  state  of  aggregation 
would  be  placed  in  this  fifth  class.  It  is  here  only  that  the  boundary 
between  physics  and  chemistry  is  at  present  difficult  to  define. 

Characteristics  of  Chemical  Change.  —  The  different  vari- 
eties of  chemical  change  can  now  be  tabulated  for  reference: 

1.  Combination  (pp.  12,  16*):  Two  (or  more)  substances  give  one 
substance. 

2.  Decomposition  (p.  17):    One  substance  gives    two   (or  more) 
substances. 

3.  Displacement   (p.   18):    One  simple  substance  and   one   com- 
pound substance  give  one  simple  substance  and  one  compound. 

4.  Double  Decomposition  (p.  20) :  Two  compound  substances  give 
two  compounds. 

5.  Internal  rearrangement  (p.  21) :  One  substance  gives  one  sub- 
stance. 

The  characteristics  of  chemical  change,  thus  encountered,  are, 
therefore,  (1)  that  the  change  can  almost  always  be  classified  under 
one  of  the  five  varieties  j  ust  mentioned,  and  (2)  that  each  of  the  sub- 
stances used  or  produced  has  a  distinct  set  of  specific  physical  proper- 
ties of  its  own,  by  means  of  which  it  can  be  recognized. 

*  See  footnote  on  p.  14. 


22  INORGANIC  CHEMISTRY 

Simple  and  Compound  Substances.  —  If  we  place  before  a 
physicist  samples  of  iron,  ferrous  sulphide,  and  sulphur,  he  will  re- 
port that  there  are  three  absolutely  distinct  substances  represented, 
because  they  show  three  different  sets  of  physical  properties.  A 
chemist,  on  the  other  hand,  while  admitting  the  accuracy  of  the  re- 
port, in  view  of  the  criterion  used  by  the  physicist,  which  indeed  he 
uses  himself  (cf.  p.  4),  will  insist  on  adding  that  there  are  only  two 
perfectly  distinct  kinds  of  -matter  in  the  set,  because  he  can  make  the 
second  from  matter  furnished  by  the  other  two.  The  same  sharp 
contrast  in  the  points  of  view  arises  when  mercury,  mercuric  oxide, 
and  oxygen,  or  any  similar  set  of  substances,  is  submitted  to  the  same 
two  tribunals.  In  a  sense,  chemistry  reduces  the  kinds  of  different 
matter  to  a  much  smaller  number  than  does  physics  or  any  of  the 
other  sciences,  and  so  it  is  the  final  authority  in  all  questions  involving 
matter.  By  the  chemist,  dozens  of  physically  distinct  substances  are 
regarded  as  closely  related  because  they  all  can  be  made  with  iron, 
or  when  decomposed  give  it;  hundreds  are  alike  in  that  sulphur 
enters  into  their  composition;  thousands  are  compounds  of  oxygen. 
In  fact,  the  number  of  kinds  of  matter  which  are  perfectly  distinct 
in  the  strictly  chemical  point  of  view  is  quite  limited. 

The  conception  contained  in  the  last  statement  was  not  reached 
until  centuries  of  effort  had  been  spent  in  trying  to  make  gold  out  of 
pyrite  (a  shining  yellow  mineral),  or  silver  out  of  lead.  The  first 
to  put  our  modern  view  into  definite  language  was  Lomonossov. 
Later,  and  independently,  it  was  stated  very  clearly  by  Lavoisier  in 
his  Traite  de  Chimie  (1789).  Lomonossov  never  believed  in  phlogis- 
ton (p.  9),  although  ail  his  contemporaries  did,  and  the  later  inves- 
tigations of  Lavoisier  finally  overthrew  this  absurd  hypothesis  and 
caused  the  general  acceptance  of  the  view  that  chemical  changes 
involved  only  combination  or  decomposition  of  different  kinds  of 
matter.  His  work  showed,  also,  that  decomposition  had  its  limits. 
Mercuric  oxide  could  be  decomposed  into  mercury  and  oxygen,  but 
no  means  was  found  of  breaking  these  up  in  turn  and  producing  any 
fresh  substances  from  them.  The  kinds  of  matter  composing  these 
simple  materials  he  named  elements.  The  element  is  to  be  regarded 
as  an  ultimate  chemical  individual  just  as  the  substance  is  the 
physical  individual.  The  definition  of  an  element  is  therefore:  a 
distinct  species  of  matter  which  we  are  not  able,  at  will,  to  decompose 
into,  or  to  make  by  chemical  union  from,  other  substances. 

The  caution  which  prompted  Lavoisier  to  use,  as  he  did,  the  words 
''has  not  yet  been/'  was  justified  by  the  fact  that  several  substances, 


CHEMICAL  PHENOMENA  23 

in  his  time  regarded  as  elementary,  were  afterwards  shown  to  be 
compound.  Thus,  quicklime  was  a  simple  substance  until  Davy,  in 
1808,  prepared  the  metal  calcium  and  showed  that  quicklime  was  a 
compound  of  this  metal  with  oxygen.  Hence,  we  do  not  say  that  the 
substances  regarded  as  simple  cannot  be  decomposed,  but  only  that 
they  are  substances  which  we  "are  not  able"  (at  present)  to  de- 
compose. 

The  phrase  "at  will"  is  also  important.  Radium  (q.v*)  cannot 
be  decomposed  at  will,  but  it  undergoes  continuous  "  disintegration," 
producing  the  elements  helium  and  lead.  We  can  neither  hasten, 
retard,  nor  stop  this  spontaneous  decomposition. 

Recent  discoveries,  showing  that  the  atoms  of  which  elementary 
substances  are  composed  are  made  up  of  different  numbers  of  par- 
ticles of  positive  and  negative  electricity,  have  revived  the  idea  that 
it  may  be  possible  to  make  one  element  out  of  another.  If  this  should 
be  accomplished,  as  it  soon  may  be,  it  may  be  necessary  radically  to 
revise  the  definition  just  given. 

The  chemist's  work  is,  at  present,  directed  wholly  by  the  thought 
that  the  individual  element  (the  matter),  after  combination,  is  still 
present  in  the  compound  in  some  form  which  is  at  least  ^item-discrete. 
The  readiness  of  the  element  to  be  released  once  more  under  suitable 
conditions  seems  to  favor  this  point  of  view. 

A  compound  is  a  substance  which  can  be  made  by  chemical  com- 
bination of,  or  can  be  decomposed  into,  two  or  more  substances. 

Elements.  —  The  word  element  is  used  in  two  senses.  It  is 
applied  to  the  simple  substance.  Thus  we  speak  of  "the  element 
iron,"  meaning  the  metal  iron.  It  is  applied  also  to  the  iron-matter 
contained  in  ferrous  sulphide  or  in  ferric  oxide.  The  reader  should 
note  that  it  is  correct  usage  to  speak  of  the  element  iron  and  the 
element  sulphur  in  ferrous  sulphide,  but  a  chemist  would  never  say 
that  this  compound  contained  the  simple  substances  iron  and  sulphur. 
If  he  did,  we  should  understand  him  to  mean  that  it  was  a  mixture, 
and  we  should  expect  parts  of  the  material  to  be  magnetic  like  iron, 
and  other  parts  to  be  yellow  and  soluble  in  carbon  disulphide,  which 
is  not  the  case.  In  the  same  way  the  name  of  an  element  (such  as 
"iron")  is  applied  both  to  the  material  in  combination  and  to  the 
free  substance.  Thus  "iron"  may  mean  free,  uncombined,  metallic 
iron,  or  iron-matter  in  some  compound.  The  sense  in  which  the 
word  is  employed  must  be.  inferred  from  the  context  or  circum- 

*  See  footnote  on  p.  14. 


24  INORGANIC  CHEMISTRY 

stances.  When  a  chemist  speaks,  as  he  sometimes  does,  colloquially, 
of  "iron"  in  a  drinking  water,  for  example,  we  know  at  once  that  he 
refers  to  iron  in  the  form  of  some  compound,  for  metallic  iron  does 
not  dissolve  in  water  and,  if  it  did,  would  quickly  turn  into  rust  or 
some  other  form  of  combination. 

The  word  element,  then,  means  one  of  the  simple  forms  of  matter, 
either  free  or  in  combination. 

In  formally  describing  a  body  or  specimen,  the  chemist  always 
avoids  the  ambiguity  just  referred  to  by  naming  the  components,  i.e., 
the  substance  or  substances  it  contains.  He  assumes  that  the  nature 
and  constituents  of  these  substances  will  be  known  to  anyone  hearing 
or  reading  the  description.  If  he  says  the  body  contains  zinc  and 
sulphur,  it  is  understood  that  the  body  is  a  mixture  of  these  simple 
substances.  If  it  contains  these  elements  in  combination,  the 
chemist  would  report  that  it  was  sulphide  of  zinc. 

The  Common  Elements.  —  Thousands  of  different  compound 
substances  are  known  but,  when  they  are  decomposed,  it  is  found 
that  the  number  of  different  elements  contained  in  them  is  not  great. 
As  we  have  said,  dozens  of  substances  contain  iron,  hundreds  contain 
sulphur,  thousands  contain  oxygen.  In  fact,  by  combining  a  limited 
number  of  simple  substances,  two,  three,  or  four  together,  in  varying 
proportions  by  weight,  an  almost  unlimited  number  of  different 
compound  substances  could  be  produced. 

The  list  of  the  elements  appears  on  the  inside  of  the  cover,  at 
the  end  of  this  book,  and  contains  about  eighty  names.  Of  these, 
a  large  number  are  rare,  and  seldom  encountered.  More  than 
99  per  cent  of  terrestrial  materials  is  made  up  of  eighteen  or  twenty 
elements  and  their  compounds.  Only  about  twenty  elements  occur 
in  nature  in  their  simple,  uncombined  condition.  Three-fourths  of 
the  whole  number  are  found  in  combination  exclusively,  and  must 
be  liberated  by  some  chemical  action. 

Taking  the  atmosphere,  all  terrestrial  waters,  and  the  earth's 
crust,  so  far  as  it  has  been  examined,  F.  W.  Clarke  has  estimated 
the  plentifulness  of  the  various  elements.  The  first  twelve,  with 
the  quantity  of  each  contained  in  one  hundred  parts  of  terrestrial 
matter,  are  as  follows: 

Oxygen 49.85  Calcium 3.18  Hydrogen 0.97 

Silicon 26.03  Sodium 2.33  Titanium 0.41 

Aluminium 7.28  Potassium 2.33  Chlorine 0.20 

Iron 4. 12  Magnesium 2.11  Carbon 0. 19 


CHEMICAL  PHENOMENA  25 

Thus  oxygen  accounts  for  nearly  one-half  of  the  whole  mass.  Silicon, 
the  oxide  of  which  when  pure  is  quartz  and  in  less  pure  form  con- 
stitutes the  ordinary  sand,  makes  up  half  of  the  remainder.  Valu- 
able and  useful  elements,  like  gold,  silver,  sulphur,  and  mercury,  are 
among  the  less  plentiful  which,  all' taken  together,  furnish  the  re- 
maining one  per  cent. 

Warnings.  —  A  substance  (p.  4)  must  be  defined  by  the  con- 
stancy of  its  physical  properties,  and  not  by  saying  that  it  is  "a 
material  of  some  specified,  definite  composition."*  We  know  many 
sets  of  anywhere  from  two  to  a  dozen  or  more  substances  with  the 
very  same  composition,  yet  the  members  of  each  set  have  entirely 
different  physical  properties,  and  often  are  entirely  different  in 
chemical  behavior.  Thus,  red  and  yellow  phosphorus  are  both 
composed  entirely  of  phosphorus,  yet  differ  markedly  in  physical 
properties.  Again,  urea  and  ammonium  cyanate  have  the  same 
composition,  but  differ  in  physical  properties  and,  besides,  belong 
chemically  to  entirely  different  classes  of  substances.  The  first  is 
basic  in  nature,  the  second  a  salt. 

Again,  a  substance  must  not  be  defined  as. "a  portion  of  matter," 
for  that  is  the  definition  of  a  body  (p.  7).  A  substance  is  a  variety 
or  species  of  matter,  and  not  any  one  portion  or  specimen.  A  crystal 
of  salt  is  "a  portion  of  matter,"  and  a  specimen  of  the  substance,  but 
the  phrase  "the  substance  salt"  covers  all  known  and  unknown 
specimens  of  salt,  and  even  not-yet-existent  specimens  which  may 
be  formed  in  the  future. 

Beware  of  defining  a  compound  substance  (p.  12)  as  being  "com- 
posed of  two  or  more  simple  substances."  That  is  a  definition  of  a 
mixture,  and  such  a  product  would  show  two  or  more  kinds  of  matter 
with  different  properties.  A  compound  substance  contains  only  one 
substance,  and  every  particle  has  the  same  physical  properties  as 
every  other  particle.  A  compound  substance  might  be  defined  as 
"a  substance  composed  of  two  or  more  elements."  This  definition, 
however,  is  academic,  and  gives  no  clew  to  how  a  compound  is  shown 
to  be  such.  Chemistry  is  an  experimental  subject,  and  all  defini- 
tions should  be  stated  in  terms  of  the  experimental  method  of  classify- 
ing a  given  case,  as  is  done  in  the  definition  given  above  (p.  12). 

*  In  chemistry,  the  word  composition,  when  applied  to  a  substance,  refers 
both  to  the  elements  present  in  the  compound  and  to  the  proportions  by  weight 
in  which  they  are  present. 


26  INORGANIC  CHEMISTRY 

Summary.  —  In  the  latter  half  of  this  chapter  we  have  learned; 
(1)  that  physical  properties  are  utilized  in  manipulations,  like  ni- 
tration and  evaporation,  as  well  as  for  identifying  substances;  (2) 
that  practically  all  chemical  changes  can  be  classified  under  one  of 
five  varieties;  (3)  that,  while  there  are  very  many  substances,  there 
is  a  very  limited  number  of  entirely  different  kinds  of  matter  (ele- 
ments). 

Exercises.*  —  1.  Take  one  by  one  the  words  or  phrases  printed 
in  black  type  and  the  titles  of  the  sections  in  this  chapter,  and 
endeavor  to  recollect  what  you  have  read  about  each.  In  each  case 
try,  (a)  to  recall  the  meaning  and  to  state  it  in  your  own  words ;  (b) 
to  recall  the  facts  associated  with,  and  the  reasoning  which  led  up  to 
the  point  in  question;  (c)  to  recall  examples  illustrating  the  concep- 
tion and  to  apply  the  conception  in  detail  to  each  example.  When- 
ever memory  fails  to  give  a  perfectly  clear  report  of  the  matter  in 
hand,  the  text  must  be  read  and  re-read  until  the  essential  point  can 
be  repeated  from  memory. 

Use  the  same  method  in  all  future  chapters.  A  useful  practice 
is  to  employ  a  pencil  as  you  read  and  to  underline  systematically  all 
the  important  facts  and  statements,  and  then  to  go  back  and  apply 
to  each  marked  place  the  process  described  above. 

2.  Define  the  following  terms:   density,  tenacity,  melting-point, 
brittleness,  specific  physical  property,  pure  body,  vacuum. 

3.  Is  it  logical  to  say  "pure  substance"? 

4.  Why  do  we  decide  that  granite  is  a  mixture  and  iron  a  single 
substance? 

5.  Do  the  statements  in  the  text  indicate  that  air  is  a  mixture  or 
a  compound? 

6.  What  weight  of  oxygen  would  be  required  to  turn  25  grams  of 
lead  into  oxide  of  lead? 

7o  Make  a  list  of  the  technical  words  we  have  defined,  and  place 
the  definitions  opposite  to  each. 

8.  What  weight  of  tin  would  be  contained  in  15  grams  of  oxide 
of  tin? 

9.  If  any  of  the  following  are  mixtures,  mention  the  facts  which 

*  The  exercises  should  in  all  cases  be  studied  with  minute  care.  They 
not  only  serve  as  tests  to  show  that  the  chapter  has  been  understood,  but  very 
frequently  also  call  attention  to  ideas  which  might  not  be  acquired  from  the  text 
alone,  or  (as  in  Nos.  11,  12)  assist  in  elucidating  ideas  given  in  the  text  which, 
without  the  exercises,  might  not  be  fully  grasped. 


CHEMICAL  PHENOMENA  27 

show  them  to  contain  more  than  one  substance:  (a)  muddy  water, 
(b)  an  egg,  (c)  milk,  (d)  a  candle,  (e)  a  cake  of  soap. 

10.  What  are  the  two  most  direct  ways  of  showing  a  substance 
to  be  a  compound?     Illustrate  each. 

11.  If  we  say  that  quicklime  contains  calcium  (p.  23),  do  we 
mean  the  element  or  the  simple  substance  calcium? 

12.  What  physical  properties  are  used,  (a)  in  nitration,  (b)  in 
evaporation,  (c)  in  the  separation  and  identification  of  the  products 
from  heating  mercuric  oxide  (p.  17)? 

13.  Take  the  chemical  action  on  p.  16,  par.  3,  and  enumerate  the 
physical  properties  of  the  substances  before  and  after  the  chemical 
change. 

14.  Discuss  in  detail  the  experiments  with  zinc  and  sulphuric 
acid  (p.  18),  and  with  silver  nitrate  (p.  19),  showing  what  specific 
properties  were  used  for  separating  and  identifying  the  products, 
and  how  they  answered  the  purpose.     Which  methods  of  manipu- 
lation were  employed  in  the  second  experiment,  and  which  method 
was  used,  essentially,  in  the  first? 

15.  Define  the  following  terms,  and  find  illustrations  of  each, 
other   than   those   given   on  p.   6:    mixture,   physical  component, 
chemical  constituent. 

16.  What  (a)  elements  and  (b)  substances  are  contained  in  an 
aqueous  solution  of  sodium  nitrate?     Would  it  be  correct  to  say 
that  the  simple  substance  oxygen  is  contained  in  it?     What  then 
is  the  difference  in  meaning  between  the  terms  "element  oxygen" 
and  the  " simple  substance  oxygen"? 

17.  What  explanation  was  given,   (a)  of  the  disappearance  of 
mercuric  oxide  when  heated,  (b)  of  the  absence  of  iron  and  sulphur, 
as  substances,  from  ferrous  sulphide?     Which  of  the  three  kinds  of 
explanation  was  used  in  each  case? 


CHAPTER  II 

ENERGY  IN  CHEMICAL  CHANGE.    PHYSICS  IN  PRACTICAL 

CHEMISTRY 

STUDY  of  the  four  typical  chemical  changes  described  in  the  last 
chapter  may  now  be  resumed,  in  order  to  see  whether  anything  further 
of  a  general  nature  is  characteristic  of  such  phenomena. 

Physical  Concomitants  of  Change  in  Composition.  —  We 

recall  at  once  that  a  prominent  feature  of  the  union  of  iron  and  sulphur 
(p.  16)  was  the  heat  which,  as  shown  by  the  glow  spreading  through 
the  mass,  seemed  to  be  developed  after  the  action  was  once  started. 
It  is  found  that  many  chemical  changes  are  like  this  one,  in  exhibiting 
simultaneously  the  production  of  very  perceptible  amounts  of  heat. 
The  burning  of  wood  and  of  coal  are  examples.  On  the  other  hand, 
the  decomposition  of  mercuric  oxide,  as  was  pointed  out  (p.  17), 
owed  its  continuance  to  the  persistent  application  of  heat,  and  ceased 
so  soon  as  the  source  of  heat  was  withdrawn.  Here,  apparently, 
heat  was  consumed  during  the  progress  of  the  change,  and  the  chemi- 
cal action  was  limited  by  the  amount  of  heat  supplied.  The  pro- 
duction or  consumption  of  heat  may,  therefore,  be  a  feature  of 
chemical  change. 

In  the  iron  and  sulphur  case,  as  in  other  chemical  actions  where 
the  heat  developed  is  great,  light  also  was  given  out.  In  the  last  of 
the  actions,  on  the  other  hand,  we  obtained  a  substance  (silver 
chloride),  which  may  be  kept  for  any  length  of  time  in  the  dark,  but, 
by  the  action  of  sunlight  is  broken  up  into  its  constituents  (p.  19). 
It  would  appear,  therefore,  that  light  may  be  given  out  or  used  in 
connection  with  chemical  change.  Noting  these  facts  stimulates  us 
to  look  for  other  similar  concomitants  of  change  in  composition. 

If  we  dip  two  wires  from  a  battery  or  dynamo  into  a  solution  of 
nitrate  of  silver  (Fig.  16),  such  as  was  used  in  the  fourth  experiment, 
we  observe  the  instant  production  of  a  coating  of  silver  on  the  nega- 
tive wire.  By  preparing  the  solution  properly  and  allowing  the 
electricity  to  flow  through  it  for  a  sufficient  length  of  time,  all  of 
the  compound  can  be  decomposed  and  all  its  silver  deposited.  It  is 

28 


ENERGY  IN  CHEMICAL  CHANGE 


29 


needless  to  say  that  this  release  of  the  silver  from  chemical  combina- 
tion, and  liberation  of  the  metal  at  the  electrode,  goes  on  only  so  long 
as  the  current  of  electricity  is  employed,  and  that  electrical  energy  is 
consumed  in  the  process.  Very  many  substances  can  be  decomposed 
in-  this  way. 

The  inverse  of  this  is  likewise  familiar.  If  we  place  in  dilute 
sulphuric  acid  a  stick  of  the  metal  zinc  (p.  18),  we  find  that  a  gas 
is  given  off  rapidly  (Fig.  17),  that  the  zinc  gradually  dissolves,  and 


FIG.  16. 


FIG.  17. 


that  a  large  amount  of  heat  is  developed.  A  thermometer  immersed 
in  the  vessel  will  show  that  the  temperature  is  rising.  If  much 
pulverized  zinc  is  used,  the  liquid  may  even  rise  spontaneously  to  the 
boiling-point.  This  form  of  the  action  produces  heat.  If,  however, 
we  attach  the  same  stick  of  zinc  to  a  copper  wire,  and  immerse  it 
and  a  plate  of  platinum  simultaneously  in  the  acid  (Fig.  18),  then  a 
galvanometer,  with  which  the  wires  are  connected,  shows  at  once 
the  passage  of  a  current  of  electricity  round  the  circuit.  Exactly 
the  same  chemical  change  goes  on  as  before  (p.  18).  The  sole 
difference  is  that  the  gas  appears  to  arise  from  the  surface  of  the 
platinum.  It  is  easy  to  show,  however,  that  the  platinum  by  itself 
is  not  acted  upon  by  dilute  acids,  and,  in  this  case,  undergoes  no 
change  whatever;  it  serves  simply  as  a  suitable  conductor  for  the 
electricity.  Here,  then,  in  place  of  the  heat  which  the  first  plan 


30 


INORGANIC  CHEMISTRY 


produced,  we  get  electrical  energy.  The  arrangement  is,  in  fact,  a 
battery-cell,  for  a  battery  is  a  system  in  which  a  chemical  action, 
which  would  otherwise  give  heat,  furnishes  electricity  instead. 
Thus,  electrical  energy  may  be  consumed  or  produced  in  connection 
with  a  change  in  composition. 


FIG.  18. 

Even  violent  rubbing  in  a  mortar,  in  the  case  of  some  substances, 
can  effect  an  appreciable  amount  of  decomposition  in  a  few  minutes. 
In  this  way  silver  chloride  can  be  separated  into  silver  and  chlorine, 
just  as  by  light.  It  is  the  mechanical  energy  which  is  the  agent,  and 
part  of  it  is  consumed  in  producing  the  change,  and  only  the  balance 
appears  as  heat.  Conversely,  the  production  of  mechanical  energy,  as 
the  result  of  chemical  change,  is  seen  in  the  behavior  of  explosives 
and  in  the  working  of  our  muscles.  Thus,  mechanical  energy  may  be 
used  up  or  produced  in  chemical  changes. 

Summing  our  experience  up,  we  may  state  that  no  change  in  com- 
position occurs  without  some  concomitant,  such  as  the  production  or 
consumption  of  heat,  light,  electrical  energy,  or,  in  some  cases, 
mechanical  energy.  It  must  be  noted  that  these  phenomena  are  an 
essential  part  of  the  chemical  change,  and  are  as  important  as  is 
the  production  of  new  substances  which  goes  on  at  the  same  time. 
We  must,  therefore,  give  attention  to  both. 

Classification  of  the  Concomitants  of  Change  in  Composi- 
tion: Energy.  —  The  problem  of  classifying  (i.e.,  placing  in  a  suit- 
able category)  things  like  heat,  light,  and  electricity  has  occupied 


ENERGY  IN  CHEMICAL  CHANGE 


31 


much  attention.  They  do  not  possess  mass.  In  all  chemical  changes, 
one  of  these  natural  concomitants  is  given  out  or  absorbed,  sometimes 
in  great  amount,  yet  in  none  is  any  alteration  in  weight  observed. 
Nor  may  these  concomitants  be  overlooked,  simply.  A  conception 
is  a  real  thing;  a  religious  belief  may  be  most  real  and  potent.  There 
are  many  things  which  are  real,  although  they  are  not  affected  by 
gravitation.  In  the  present  instance  we  reason  as  follows: 

A  brick  in  motion  is  different  from  a  brick  at  rest.  The  former 
can  do  some  things  that  the  latter  cannot.  Furthermore,  we  can 
easily  make  a  distinction  in  our  minds.  The  brick  can  be  deprived 
of  the  motion  and  be  endowed  with  it  again.  Thus,  we  can  get  the 
idea  of  motion  as  a  separate  conception.  Similarly,  we  observe  that 
a  piece  of  iron  behaves  differently  when  hot,  and  when  cold,  when 
bearing  a  current  of  electricity,  and  \\hen  bearing  none.  We  con- 
ceive then  of  the  brick  or  the  iron  as  having  a  certain  amount  and 
kind  of  matter  which  is  unalterable,  and  as  having  motion,  heat,  or 
electricity  added  to  this  or  removed.  Thus,  we  describe  our  observa- 
tions by  using  two  categories,  one  of  which 
includes  the  various  kinds  of  matter,  and 
the  other,  various  things  whose  association 
with  matter  seems  to  be  invariable  and  is 
often  so  conspicuous. 

At  first  sight,  these  concomitants  of 
matter  seem  to  be  quite  disparate.  But 
a  relation  between  them  can  be  found. 
If  the  heat  of  a  Bunsen  flame  or  of  the 
sun  is  brought  under  a  hot-air  motor  (Fig. 
19)  violent  motion  results.  Again,  if  the 
motor  is  connected  with  a  dynamo,  elec- 
tricity may  be  generated.  Still  again,  if 
the  current  flows  through  an  incandescent 
lamp,  heat  and  light  are  evolved.  Con- 
versely, when  motion  is  impeded  by  a 
brake,  heat  appears.  When  a  current  of  electricity  is  run  through 
the  dynamo,  motion  results.  But  the  most  significant  facts  are  still  to 
be  mentioned.  The  heat  absorbed  by  the  motor  is  found  to  be  greater 
when  the  machine  is  permitted  to  move  and  do  work,  than  when 
it  is  not.  Thus,  we  find  that  when  work  is  done  by  the  motor  some 
heat  disappears,  being  transformed  into  work.  Similarly,  when  the 
poles  of  the  dynamo  are  properly  connected  and  electricity  is  being 
produced,  and  only  then,  motion  is  used  up.  This  is  shown  by  the 


FIG.  19. 


32  INORGANIC  CHEMISTRY 

effort  required  to  turn  the  armature  under  these  circumstances,  and 
the  ease  with  which  it  is  turned  when  the  circuit  is  open.  So,  with  a 
conductor  like  the  filament  in  the  lamp,  unless  it  offers  resistance 
to  the  current  and  destroys  a  sufficient  amount  of  electricity,  it 
gives  out  neither  light  nor  heat.  Finally,  motion  gives  no  heat 
unless  the  brake  is  set,  and  effort  is  then  demanded  to  maintain  the 
motion.  These  experiences  lead  us  to  believe  that  we  have  here  a 
set  of  things  which  are  fundamentally  of  the  same  kind,  for  each 
form  can  be  made  from  any  of  the  others.  We  have,  therefore, 
invented  the  conception  of  a  single  thing  of  which  heat,  light,  elec- 
tricity, and  motion  are  forms,  and  to  it  we  give  the  name  energy: 
energy  is  work  and  every  other  thing  which  can  arise  from  work  and  be 
converted  into  work  (Ostwald). 

Closer  study  shows  that  equal  amounts  of  electrical  or  mechanical 
energy  always  produce  equal  amounts  of  heat.  There  is  never  ob- 
served any  loss  in  transformations  of  energy,  any  more  than  in 
transformations  of  matter.  Hence,  J.  R.  Mayer  (1842),  Colding 
(1843),  and  Helmholtz  (1847)  were  led  independently  to  the  con- 
clusion that  in  a  limited  system  no  gain  or  loss  of  energy  is  ever 
observed.  This  brief  statement  of  the  results  of  many  experiments 
is  called  the  law  of  the  conservation  of  energy. 

A  current  form  of  this  law,  namely,  that  "  the  total  amount  of  energy  in  the 
universe  is  a  constant  quantity,"  is  open  to  the  same  objection  as  the  correspond- 
ingly flamboyant  form  of  the  law  of  conservation  of  mass  criticized  later  (p.  52). 
It  has  a  more  effective  sound  than  the  one  we  have  given.  Unfortunately,  it 
is  not  only  immensely  in  excess  of  any  statement  that  present  results  of  scientific 
work  can  justify,  but  is  probably  far  beyond  the  limits  of  possible  scientific 
observation.  Scientific  statements  of  fact  can  never  err  by  being  too  conservative. 

Matter  and  Energy  as  Concepts,  and  Definitions  of  the 
Latter.  —  The  foregoing  paragraphs  about  energy  bring  up  the  question  of  its 
relation  to  matter.  This  relation  can  be  made  clear  only  by  a  somewhat  elaborate 
discussion  of  our  fundamental  conceptions. 

The  only  real,  first-hand  knowledge  which  we  possess,  is  that  of  our  states  of 
consciousness.  All  else,  consisting,  for  example,  of  the  way  in  which  we  inter- 
pret and  describe  our  experience,  is  constructed  out  of  our  heads,  so  to  speak. 
Now,  we  become  aware  of  certain  things  which  we  call  sensations,  and  seek  to  con- 
struct a  mode  of  correlating  and  describing  them.  Our  universal  habit  is  to  speak 
as  if  they  were  produced  by  something  outside  our  minds,  and  so  we  begin  the 
manufacture  of  an  external  universe.  In  course  of  doing  this,  we  encounter 
some  things  which  seem  to  occupy  no  particular  space,  which  move  from  object 
to  object,  and  possess  no  weight.  One  of  these  affects  our  eye,  or  a  piece  of 


ENERGY  IN  CHEMICAL  CHANGE  33 

chloride  of  silver,  for  example,  yet  escapes  touch,  and  passes  through  glass  almost 
as  easily  as  through  a  vacuum.  After  consideration  of  our  experience  with  this 
sort  of  thing,  some  of  which  has  been  detailed  above,  we  decide  that  we  shall  posit 
the  existence  of  energy. 

Other  things  we  encounter  appeal  to  the  sense  of  touch  and  seem  to  possess 
more  definitely  located  qualities,  including  weight.  Another  conception  is  needed 
to  account  for  these;  so  we  establish  the  category  of  matter. 

Thus,  we  make  shift  to  describe  our  sensations  by  the  help  of  these  two  con- 
structs, much  as  in  analytical  geometry  we  describe  the  location  of  a  point  by 
means  of  two  coordinates.  Energy  and  matter  are,  therefore,  products  of  thought 
and  not,  primarily,  objective  realities.  In  chemistry,  however,  we  always  speak 
of  them  objectively.  Historically,  the  order  in  which  these  two  concepts  were 
named  and  denned  was  the  opposite  of  that  in  which  they  stand  above.  Yet 
attempts  to  organize  a  conception,  corresponding  to  energy,  in  response  to  a  need 
of  which  thinkers  were  conscious,  were  not  wanting  before  the  nineteenth  century 
opened.  To  go  no  further  back  than  the  days  of  phlogiston,  we  can  easily  per- 
ceive a  certain  resemblance  between  this  concept  and  that  of  energy.  The  idea 
that  heat  was  an  "imponderable"  had  its  origin  much  earlier,  and  shows  the  ex- 
istence of  the  same  effort  to  find  a  second  fundamental  conception  different  from 
matter. 

There  is  much  confusion  of  thought  in  many  of  the  current  definitions  of 
energy.  For  example,  it  is  often  said  to  be  "that  which  causes  change  in  matter." 
This  definition  is  not  easy  to  bring  into  harmony  with  common  experience  in 
chemistry.  Thug,  when  heat  is  applied  to  mercuric  oxide,  the  change  follows. 
But  with  iron  and  sulphur,  the  union  of  the  two  substances  is  a  condition  antece- 
dent to  the  evolution  of  heat.  It  is  as  often  true  that  change  in  matter  causes  the 
manifestation  of  energy  as  the  reverse.  Matter  and  energy  are  on  the  same  plane. 
They  are  conceptions  used  jointly  in  describing  what  we  observe.  Neither  is 
secondary  to  the  other.  We  do  not  consider  any  particular  one  of  the  coordinates 
in  geometry  as  secondary  to  the  other,  or  as  being  affected  by  the  other.  • 

The  theory  of  chemical  potential  and  of  the  factors  of  energy  (q.v.}  seeks 
once  more  to  ascribe  the  tendency  to  change  (physical  or  chemical)  in  matter  to 
the  state  of  the  energy  associated  with  it.  It  is,  therefore,  incidentally,  so  con- 
structed as  to  favor  the  definition  just  given. 

The  definition  that  "matter  is  the  vehicle  of  energy"  is  obviously  just  as  diffi- 
cult to  harmonize  with  the  above  mode  of  deriving  the  two  conceptions.  One  axis 
is  not  spoken  of  as  the  vehicle  of  the  other  in  geometry. 

The  innate  desire  to  reduce  our  distinct  categories  to  the  smallest  possible 
number  may  be  seen  in  the  history  of  this  subject.  The  ancients  sought  the 
amalgamation  of  the  two  by  regarding  heat  and  light  as  imponderable  forms  of 
matter.  It  is  now  believed  that  the  one  conception  of  energy  is  sufficient,  and 
that  matter  may  be  put  into  the  same  category  as  being  composed  of  minute 
particles  of  electricity  (positive  charges  and  electrons). 

The  conception  of  ether  was  devised  because  those  of  matter  and  energy  did 
not  suffice  for  the  description  of  all  the  phenomena  of  light.  It  is  on  the  same 


34  INORGANIC  CHEMISTRY 

plane  with  matter  and  energy.  Lord  Kelvin's  effort  to  reduce  the  three  catego- 
ries to  two  (energy  and  ether)  by  assigning  the  r61e  of  matter  to  vortices  in  the 
ether  is  familiar  to  students  of  physics. 

Application  of  the  Conception  of  Energy  in  Chemistry.  — 

At  first  sight  it  looks  as  if  the  statement  that  energy  is  conserved  is 
not  applicable  in  chemistry.  Heat  and  electricity,  for  example,  seem 
to  be  produced  and  consumed,  in  connection  with  changes  in  composi- 
tion, in  a  mysterious  manner.  We  trace  light  in  an  incandescent 
lamp  back  to  the  electricity,  and  this  in  turn  to  the  mechanical  energy, 
and  this  again  to  the  heat  in  the  engine.  But  what  form  of  energy 
gave  the  heat  developed  by  the  combustion  of  the  coal  under  the 
boiler,  or  by  the  union  of  iron  and  sulphur  in  our  illustration  (p.  16)? 
Since  we  do  not  perceive  any  electricity,  light,  heat,  or  motion  in 
the  original  materials,  and  yet  wish  to  create  an  harmonious  system, 
we  are  bound  to  conceive  of  the  iron  and  the  sulphur,  and  the  coal 
and  the  air,  as  containing  another  form  of  energy,  which  we  call 
chemical  or  internal  energy.  Similarly,  when  heat  is  used  up  in 
decomposing  mercuric  oxide,  or  light  in  decomposing  silver  chloride, 
we  regard  the  energy  as  being  stored  in  the  products  of  decom- 
position in  the  form  of  internal  energy. 

The  Actual  Quantities  of  Different  Kinds  of  Energy  which 
may  he  Obtained  from  a  Fixed  Amount  of  One  Kind.  —  It  will 
render  all  the  above  clearer  if  we  give  some  numerical  illustrations:  A  kilogram 
of  water  after  falling  (hi  a  vacuum)  428  meters  (about  one-fifth  of  a  mile),  under 
gravity,  possesses  428  kilogram-meters  of  mechanical  (kinetic)  energy.  When  the 
motion  is  arrested,  the  energy  of  motion  is  transformed  into  heat  and  raises  the 
water  one  degree  centigrade  in  temperature.  We  describe  this  amount  of  heat 
as  1000  calories  (small);  that  required  to  warm  one  gram  of  water  one  degree 
(at  15°)  being  called  one  calorie.  A  kilogram  of  any  other  falling  material  would 
give  the  same  amount  of  heat  (1000  cal.),  although,  of  course,  if  its  specific 
heat  were  smaller  than  that  of  water,  the  temperature  to  which  it  would  be  raised 
would  be  higher,  and  vice  versa. 

Here  the  acting  force  is  the  attraction  of  gravitation,  which  is  a  special  case. 
In  absolute  units,  1  g.  falling  1  cm.  generates  energy  enough  to  do  981  ergs  of 
work.  So  that  the  thousand  grams  falling  428  meters  (42,800  cm.)  generates 

1000  X  42,800  X  981  =  42,000,000,000  ergs  of  energy. 

The  erg  being  so  small,  we  often  use  the  joule  ( =  10,000,000  ergs).     This  amount 
is  the  same  as  4200  joules. 

Now,  any  body  of  the  same  mass,  moving  with  the  same  final  velocity,  how- 
ever set  in  motion,  will  also  possess  the  same  energy  and  give  1000  cal.  The  final 


ENERGY  IN  CHEMICAL  CHANGE  35 

velocity  in  the  above  case  is  V  2  gs.  =  9164  cm.  per  second.  The  energy  of 
motion  Q  raw2)  of  one  thousand  grams  of  matter  moving  with  this  velocity  is 
=  £  X  1000  X  (9164)2  =  4200  joules,  as  before.  If  the  source  of  mechanical 
energy  were  a  hot-air  motor  (or  an  engine)  of  one  horse-power,  then,  since  one 
horse-power  represents  a  development  of  746  joules  per  second,  the  4200  joules  of 
energy  would  be  produced  in  about  5£  seconds  by  this  means. 

If, 'instead  of  being  turned  into  heat,  all  the  energy  of  motion  had  been  con- 
verted into  electricity,  the  quantity  of  the  latter  would  have  illuminated  a  50- 
watt  incandescent  lamp  for  84  seconds  (  =  1.4  minutes).  Such  a  lamp  requires 
50  joules  per  second  of  electricity  and,  therefore,  in  84  seconds  uses  up  the 
50  X  84  =  4200  joules  of  energy.  As  an  engine  of  1  horse-power  produces  this 
amount  of  energy  every  5?  seconds,  such  an  engine,  if  none  of  the  energy  were 
lost,  could  maintain  nearly  15  lamps  of  this  kind. 

Finally,  if  the  4200  joules  of  electrical  energy  were  applied  to  decomposing 
nitrate  of  silver  in  ordinary  aqueous  solution,  it  would  liberate  6f  grams  (about 
i  oz.)  of  silver  from  combination. 

Considerations  Connected  with  Internal  Energy:  Free 
Energy.  —  These  conclusions  compel  us,  for  the  sake  of  consistency, 
to  think  of  all  our  materials  as  repositories  of  energy  as  well  as  of 
matter,  each  of  these  constituents  being  equally  real  and  equally 
important.  A  piece  of  the  substance  known  as  "iron"  must  thus  be 
held  to  contain  so  much  iron  matter  and  so  much  internal  energy. 
So  ferrous  sulphide  contains  sul  phur  matter,  iron  matter,  and  internal 
energy.  Thus,  by  a  substance  we  mean  a  distinct  species  of  matter, 
simple  or  compound,  with  its  appropriate  proportion  of  internal 
energy.  During  the  progress  of  a  chemical  change,  like  the  union 
of  iron  and  sulphur,  the  chemical  energy  of  the  system  diminishes  and 
heat  is  liberated,  or,  when  arrangements  are  made  for  utilizing  the 
energy,  work  of  some  kind  is  done . 

The  energy  which  becomes  available  as  the  result  of  a  chemical  action,  and 
is  free  to  be  converted,  say,  into  electrical  energy,  is  called  the  free  energy  of 
the  action.  Now,  it  must  be  noted  that  the  free  energy,  measured  by  work  done, 
is  not,  in  general,  the  equivalent  of  the  heat  developed  by  the  natural  progress  of 
the  change.  Often  the  amounts  are  nearly  equivalent,  although  never  absolutely 
so.  But  frequently  they  are  very  different.  When  the  free  energy  available  for 
conversion  into  work  is  greater  in  amount  than  the  heat  of  the  reaction,  as  it  often 
is,  the  difference  is  taken  up  from  the  heat  of  the  surroundings  during  the  progress 
of  the  change,  and  the  vessels  and  objects  in  contact  with  the  interacting  bodies 
become  colder.  Thus  phosphonium  chloride  (q.v.)  decomposes  spontaneously 
into  two  gases,  phosphine  and  hydrogen  chloride,  and  ammonium  carbonate  gives 
off  ammonia  gas,  while  heat  is  absorbed  in  both  cases.  Work  can  be  done  by 
both  these  actions,  although,  so  far  as  heat  is  concerned,  not  only  is  none  of  this 


36  INORGANIC  CHEMISTRY 

form  of  energy  liberated,  but  a  certain  amount  of  it  is  absorbed.  Conversely, 
when  the  heat  of  reaction  is  greater  than  the  equivalent  of  the  free  energy,  then, 
along  with  the  energy  which  would  be  used  to  do  work  (for  example,  by  employing 
the  action  as  a  source  of  electricity),  a  certain  amount  of  heat  which  cannot  be 
transformed  into  work  will  be  given  to  the  surroundings.  It  thus  appears  that  the 
substances  which  we  handle  are  not  only  repositories  of  energy,  but,  when  brought 
together,  also  play  the  part  of  machines  for  transforming  energy  which  they  take 
from  or  give  to  the  surroundings. 

The   Third   Characteristic  of  Chemical   Change.  —  In  the 

course  of  this  discussion  it  has  become  clear  that  it  is  a  characteristic 
of  a  chemical  change  that,  besides  a  change  in  that  state  of  the 
matter,  there  is  always  an  alteration  in  the  amount  of  internal  energy 
in  the  system.  This  alteration  involves  the  production  of  internal 
energy  from,  or  the  transformation  of  internal  energy  into,  some  other 
form  of  energy. 

Exothermal  and  Endothermal  Changes.  —  The  energy  lib- 
erated in  a  chemical  action  appears  most  commonly  in  the  form  of 
heat.  Changes  which,  like  the  union  of  iron  and  sulphur  (p.  16),  are 
accompanied  by  the  liberation  of  heat  are  called  exothermal  actions. 
Those  in  connection  with  which  heat  is  absorbed,  like  the  decom- 
position of  mercuric  oxide  (p.  17),  are  known  as  endothermal. 

It  should  be  noted  here  that  neither  the  production  nor  the  ab- 
sorption of  heat  is  an  exclusive  mark  of  chemical  change.  Physical 
changes  are  all  likewise  accompanied  by  the  liberation  or  consumption 
of  energy.  Thus,  water,  in  evaporating,  absorbs  heat,  and  liquids 
on  solidifying,  or  often  even  when  simply  mixed  with  other  liquids, 
give  out  heat. 

Practical  Importance  of  Energy  in  Chemical  Change.  — 

The  absorption  or  liberation  of  energy  accompanying  a  chemical 
transformation  of  matter  is  often,  of  the  two,  the  more  important  fea- 
ture. We  do  not  burn  coal  in  order  to  manufacture  carbon  dioxide 
gas.  We  are  glad  to  get  rid  of  the  material  product  through  the 
chimney.  It  is  the  heat  we  want.  We  do  not  employ  zinc  in  bat- 
teries with  the  object  of  making  zinc  chloride  or  zinc  sulphate.  So  we 
use  the  electrical  energy,  and  throw  the  material  products  away 
when  we  refill  the  battery  jars.  It  is  the  same  with  burning  illumi- 
nating gas  or  magnesium  powder  when  we  want  light,  and  with  eating 
food,  which  we  do,  chiefly,  to  get  energy  to  sustain  our  activity.  We 
do  not  run  electricity  for  hours  into  a  storage  battery  in  order  to  make 


ENERGY  IN  CHEMICAL  CHANGE  37 

a  particular  compound  (lead  dioxide,  for  example),  but  in  order  to 
save  and  store  the  energy  for  future  use.  In  industry  and  life,  fully 
half  the  total  amount  of  chemical  change  involved,  is  set  in  motion 
by  us,  solely  on  account  of  the  energy  changes  it  involves. 

As  will  be  seen  in  the  following  section,  observation  of  the  amount 
of  the  energy  absorbed  or  liberated  in  chemical  changes  is  also  of 
the  greatest  importance  in  the  scientific  study  of  chemical  phenomena. 

Chemical  Activity.  —  Other  things  being  equal,  actions  in  which 
there  is  a  relatively  large  loss  of  internal  energy  and,  therefore, 
usually,  a  considerable  liberation  of  heat  or  electrical  energy,  proceed 
rapidly;  that  is  to  say,  in  them  a  large  proportion  of  the  material  is 
changed  in  the  unit  of  time.  Those  in  which  less  free  energy  is 
transformed  proceed,  in  general,  more  slowly.  The  speed  of  the 
chemical  change,  and  the  quantity  of  energy  available  because  of  it, 
are  closely  related.  Now,  we  are  accustomed  to  speak  of  materials 
which,  like  iron  and  sulphur,  interact  rapidly  and  with  liberation  of 
much  energy  as  "  chemically  active."  Thus,  relative  chemical  activ- 
ity may  be  estimated: 

1.  By  observing  the  speed  of  a  change   (see  Speed  of  chemical 
actions),  or,  in  many  cases, 

2.  By  measuring  the  heat  developed  in  the  course  of  the  action 
(see  Thermochemistry),  or, 

3.  By  ascertaining  the   electromotive  force   of  the  current  the 
change  gives,  when  arranged  in  the  form  of  a  battery-cell  (see  Electro- 
motive chemistry). 

These  different  methods  will  be  discussed  in  later  sections.  It 
sriould  be  noted  here,  however,  that  the  speed  of  a  given  action  may 
be  enormously  affected  by  conditions  (see,  for  example,  Catalysis), 
and  that,  therefore,  great  caution  is  required  in  inferring  relative 
activities  from  observed  differences  in  the  speeds  of  several  actions. 
Thermal  measurements  are  also  often  misleading.  This  is  evident 
r'rom  the  fact  that  an  action  may  be  able  to  do  work,  even  although 
heat  energy  is  absorbed  during  its  progress.  Thus,  phosphonium 
chloride  (q.v.)  decomposes  of  its  own  accord  into  phosphine  and 
hydrogen  chloride,  although  it  absorbs  heat  from  surrounding  ob- 
jects in  doing  so.  The  electrical  method  of  measuring  the  available 
(free)  energy,  and  therefore  the  true  affinity,  is  in  general  the  most 
trustworthy. 

It  is  evident  that  the  chemical  activity  of  a  given  substance  will 
not  be  the  same  towards  all  others.  Thus,  iron  unites  much  more 


38  INORGANIC  CHEMISTRY 

vigorously  with  chlorine  than  with  sulphur  and,  with  identical 
amounts  of  iron,  more  heat  is  liberated  in  the  former  case  than  in  the 
latter.  With  silver,  sodium,  and  many  other  substances,  iron  does 
not  unite  at  all.  One  of  the  tasks  of  the  chemist  is  to  make  such 
comparisons  as  this  (see  Specific  chemical  properties,  p.  86).  Evi- 
dently, the  substances  containing  the  most  internal  energy  will  be 
in  general  the  most  active. 

The  6t  Cause  "  of  Chemical  Activity.  —  The  reader  will  un- 
doubtedly be  inclined  to  inquire  whether  we  can  assign  any  cause  for 
the  tendency  which  substances  have  to  undergo  chemical  change. 
Why  do  iron  and  sulphur  unite  to  form  ferrous  sulphide,  while  other 
pairs  of  elements  taken  at  random  will  frequently  be  found  to  have 
no  effect  upon  one  another  under  any  circumstances?  This  question 
is  so  likely  to  occur  to  the  reader  that  it  should  be  dealt  with  a.t  once. 
The  answer  is  that  we  do  not  know.  Questions  like  this  have  to  go 
without  answer  in  all  sciences.  What  is  the  cause  of  gravitation? 
We  know  the  facts  which  are  associated  with  the  word  —  the  fact 
that  bodies  fall  towards  the  earth,  for  example  —  but  why  they  fall 
we  are  unable  to  say.  So,  with  chemical  change,  we  can  state  all  the 
facts  we  know  about  it,  that  is,  we  can  tell  how  the  change  takes 
place,  but  even  then  we  cannot  say  why  it  takes  place. 

The  words  "affinity"  and  "attraction"  are  sometimes  advanced  as  if  they 
supplied  some  explanation  of  chemical  activity.  Now,  we  have  seen  that  an 
explanation  in  science  (p.  10)  is  a  description  of  the  details  of  some  process,  either 
in  terms  of  known  facts,  or  by  the  use  of  some  imaginary  but  plausible  and  help- 
ful machinery.  Here  no  facts  are  known.  Even  imaginary  machinery  has  not 
yet  been  conceived  by  any  one.  So  that  these  terms  are  words  simply,  and  do 
not  meet  either  of  the  conditions  required  of  an  explanation.  They  are  names 
for  "the  tendency  to  undergo  chemical  change,"  and  that  is  all. 

All  nouns,  such  as  table  or  book,  are  general  terms  applicable  to  many  more 
or  less  various  individuals.  Some  special  nouns  are  used  in  chemistry.  For  ex- 
ample, affinity  names  the  tendency  to  undergo  chemical  change,  and  distinguishes 
this  tendency  by  name  from  cohesion,  or  the  tendency  to  unite  physically.  Ca- 
talysis names  a  kind  of  chemical  change  in  which  some  specific  substance  must 
be  present,  and  influences  the  other  substances  by  contact  with  them,  yet  itself 
undergoes  no  change.  Dissociation  names  the  kind  of  chemical  change  in  which 
decomposition  occurs  with  rise  in  temperature,  and  recombination  when  the 
temperature  falls.  But  none  of  these  terms,  as  such,  is  an  explanation.  It  does 
not  explain  the  concussion  of  two  railway  trains  to  name  it  a  collision. 

Of  course,  if  we  have  some  genuine  explanation,  applicable  to  all  the  other 
known  cases  of  a  class,  any  newly  discovered  example  falls  heir  at  once  to  this 


ENERGY  IN  CHEMICAL  CHANGE  39 

explanation.  This  would  be  true  of  a  dissociation,  where  the  kinetic  theory  and 
the  law  of  mass  action  describe  the  details  of  all  such  phenomena.  Here  the  ex- 
planation lies,  not  in  the  name,  but  in  the  knowledge  we  have  of  other  instances 
of  the  same  behavior.  The  name,  of  course,  suggests  the  whole  theory,  if  such  a 
theory  exists.  But  with  affinity,  or  the  tendency  to  enter  into  chemical  action,  we 
have  no  theory  for  any  of  the  samples  of  the  class.  We  are  entirely  ignorant  as 
yet  of  the  details  of  its  mode  of  operation,  equally  so  in  every  case,  and,  in  fact, 
know  nothing  at  all  about  it  save  that  affinity  exists  and  that  we  can  measure  its 
intensity.  So  the  name  cannot  remind  us  of  any  explanation,  for  none  has  been 
suggested. 

As  words,  the  best  one  can  say  of  them  is  that  they  are  rather  unfortunately 
chosen.  Affinity  suggests  kinship,  sympathy,  or  affection.  But  the  suggestion 
that  such  human  emotions  control  the  behavior  of  iron  and  sulphur  is  too  wild 
and  too  remote  from  common  sense  to  furnish  any  assistance.  Attraction  hints 
at  some  preexisting  bond  of  a  material  kind  which  draws  the  substances  together, 
for  we  cannot  conceive  of  action  at  a  distance  without  some  intervening  medium 
of  communication.  But  we  have  no  other  evidence  of  the  existence  of  an  instan- 
taneously adjustable  harness  capable  of  drawing  materials  into  chemical  action. 
It  is  harder  to  reduce  this  idea  to  comprehensible  shape  than  to  do  without  it. 

If  we  are  still  inclined  to  think  that  these  are  more  than  class-words,  and 
do  suggest  some  explanation,  we  have  only  to  carry  the  same  idea  further  to 
be  landed  in  absurdity.  Using  similarly  crude  analogies,  we  might  suppose  that 
the  elements  were  guided  by  scent,  like  dogs,  or  by  sight,  like  birds,  or  by  feeling, 
like  fish,  and  so  on  ad  infinitum,  and  forget  that  the  fact  itself  was  after  all  much 
simpler  than  the  explanation.  Affinity  is  simply  a  fanciful  name  for  a  real  thing. 

66  Cause"  in  Science.  —  The  word  " cause"  was  employed  in 
the  heading  of  the  last  section,  and  it  will  be  observed  that  no  cause 
was  found.  This  is  the  invariable  rule  in  physical  or  chemical 
phenomena.  We  know  of  no  causes,  in  the  sense  in  which  the  word 
is  commonly  employed. 

The  word  cause  has  only  one  definite  use  in  science.  When  we 
find  that  thorough  incorporation  of  the  three  materials  is  needed  to 
secure  good  gunpowder,  we  say  that  the  intimate  mixing  is  a  cause  of 
its  being  highly  explosive.  By  this  we  simply  mean  that  intimate 
mixture  is  a  necessary  antecedent  of  the  result.  A  cause  is  a  condition 
or  occurrence  which  always  precedes  another  condition  or  occurrence. 

Misuse  of  the  word  "cause"  is  frequent.  The  law  of  gravitation  is  not  the 
cause  of  the  behavior  of  falling  bodies.  It  is  simply  a  condensed  narration  of  the 
facts  about  falling  bodies,  and  was  made  long  after  the  first  bodies  fell.  Affinity, 
or  the  tendency  to  interact  chemically,  is  the  imagined  antecedent  of  chemical 
change.  Such  causes,  if  we  call  them  causes  at  all,  are  invented  by  way  of  supply- 
ing antecedents  to  things  that  appear  to  lack  them,  that  our  sense  of  symmetry 
may  be  satisfied  withal.  They  are  occasionally  useful.  But  the  less  fiction  we 


40  INORGANIC   CHEMISTRY 

employ  in  the  science,  the  less  will  be  the  danger  that  the  student  will  mistake 
fictions  for  facts,  or  even  fall  under  the  delusion  that  it  is  a  habit  of  science  to 
spend  more  thought  in  making  gratuitous  assumptions  than  in  ascertaining  facts. 

PHYSICS  IN  PRACTICAL  CHEMISTRY 

List  of  Important  Specific  Physical  Properties.  —  We  have 
seen  that,  to  the  chemist,  knowing  the  physical  properties  of  all  sub- 
stances is  very  important.  By  means  of  the  properties,  he  recognizes 
and  describes  all  the  bodies  he  studies.  It  may  be  well,  therefore, 
here  to  give  a  list  of  the  more  important  properties,  most  of  which 
have  been  mentioned  in  connection  with  the  illustrations  we  have 
used. 

In  the  case  of  solids,  the  chief  physical  properties  the  chemist 
uses  are:  color,  odor,  crystalline  form,  hardness  (Appendix  II), 
solubility  or  non-solubility  in  water  and  occasionally  other  liquids, 
the  temperature  at  which  the  substance  melts  (melting-point),  the 
density,  and  the  conductivity  for  electricity. 

In  the  case  of  liquids,  he  notes  the  temperature  at  which  the 
liquid  boils  (boiling-point),  the  density,  the  mobility,  the  odor,  and 
the  color. 

Finally,  in  the  case  of  gases,  the  properties  commonly  mentioned 
are  the  color,  taste,  and  odor,  the  density,  the  solubility  in  water, 
and  the  ease  or  difficulty  with  which  the  gas  can  be  liquefied  (see 
critical  temperature). 

For  example,  sulphur  is  yellow,  has  little  odor,  crystallizes  in  the 
rhombic  system,  has  a  hardness  of  2.5  on  a  scale  of  ten,  has  the 
m.-p.  112.8°  C.,  is  not  perceptibly  soluble  in  water  but  dissolves  in 
carbon  disulphide  (41  parts  :  100  at  18°),  has  the  b.-p.  444.7°  C., 
the  sp.  gr.  2,  and  is  a  very  poor  conductor. 

Slight  variations  from  the  standard  properties  of  the  substance  usually  indi- 
cate the  presence  of  an  impurity  homogeneously  incorporated.  The  precise  ways 
in  which  properties  are  affected  in  such  cases  will  be  noted  under  solutions. 

The  "substance"  will  be  seen  to  be  of  an  abstract  nature.  It  is  a  conception 
built  up  by  selecting  (or  abstracting)  the  properties  common  to  all  specimens. 
Hence  we  classified  chemistry  as  an  abstract-concrete  science  (p.  2).  The  bodies 
under  observation  are  concrete,  the  classification  of  the  results  is  under  conceptions 
of  an  abstract  nature  like  this  one. 

Attributes  and  Conditions.  —  There  are  other  qualities  which 
a  body  may  possess  that  we  are  likely  to  confuse  with  the  specific 
properties.  Thus,  the  weight  of  a  piece  of  sulphur  is  not  a  property 


ENERGY  IN  CHEMICAL  CHANGE  41 

of  sulphur.  A  hundred  pieces  of  as  many  different  substances  might 
all  have  the  same  weight,  so  that  a  particular  weight  (say  2  grams) 
is  not  a  property  of  any  one  species  of  matter.  Weight,  dimensions, 
and  volume  are  attributes  of  a  body.  They  have  different  values 
for  different  bodies,  even  when  those  bodies  are  all  composed  of  the 
same  substance.  The  attributes  are  physical  in  nature.  They  are 
of  great  importance  in  chemistry,  however,  because  they  afford  the 
only  means  we  have  of  measuring  quantities  of  matter. 

There  are  still  other  qualities  which  a  body  (or  specimen  of 
matter)  may  possess.  It  has,  for  example,  a  certain  temperature, 
pressure  (state  of  compression),  motion,  kind  of  illumination,  or 
electric  charge,  and  it  may  be  in  solution  in  some  liquid.  A  body 
may  change  in  temperature,  pressure,  or  state  of  electrification,  or 
it  may  be  dissolved  in  water,  or  be  recovered  by  evaporation  of  the 
liquid,  and  yet  be  the  same  specimen.  A  hundred  specimens  of  as 
many  different  substances  may  all  have  the  same  temperature  — 
this  is  not  a  specific  property.  These  are  spoken  of  as  conditions. 
They  are  physical  conditions.  In  chemistry,  conditions  are  often 
altered  in  order  to  bring  about,  or  to  stop  chemical  change,  or  to 
modify  the  speed  with  which  it  takes  place.  Thus  we  heated  the 
lead  (raised  its  temperature)  in  order  to  hasten  the  process  of  rusting. 
If  a  substance,  or  mixture,  is  capable  of  undergoing  chemical  change, 
then  the  change  is  always  hastened  by  raising  the  temperature,  and 
is  always  delayed  or  prevented  by  lowering  the  temperature.  Simi- 
larly, changing  the  pressure  of  gas,  or  dissolving  a  substance  in  some 
liquid,  frequently  hastens  or  delays  a  chemical  change  in  which  the 
substance  takes  part.  Again,  silver  chloride  decomposed  when 
illuminated,  but  not  in  the  dark.  The  proper  physical  conditions 
are,  therefore,  considered  in  connection  with  every  chemical  operation. 
Conditions  are  used  to  produce,  modify,  or  control  chemical  change. 

Since  the  density  of  a  substance  changes  with  the  temperature, 
we  must  define  the  density  as  the  weight  of  1  c.c.  at  4°  C.,  or  some 
other  temperature.  We  must  mention  also  the  pressure,  particularly 
in  the  case  of  gases,  and  even  in  the  cases  of  liquids  and  solids  if  it 
differs  from  one  atmosphere.  Again,  the  melting-point  is  the  tem- 
perature at  which  the  substance  melts  at  one  atmosphere  pressure. 
Thus,  the  conditions  are  used  also  in  denning  the  specific  physical 
properties. 

When  the  word  "properties"  is  used  in  speaking  of  a  substance,  we  always 
refer  to  the  specific  properties  only,  for  a  substance  is  by  definition  the  bearer  of 
nothing  but  constant,  that  is,  unchangeable  qualities.  The  weight  and  volume, 


42  INORGANIC  CHEMISTRY 

the  temperature  and  pressure,  are  variable,  and  they  do  not  enter  into  the  con- 
ception of  the  substance.  This  distinction  sets  the  abstract  nature  of  the  "sub- 
stance" in  high  relief. 

Methods  of  Work  and  Observation  in  Chemistry.  —  It  is 

not  the  end  of  chemical  work  to  make  generalizations  or  laws  (pp.  6, 11, 
etc.),  or  conceptions,  like  those  dealt  with  in  the  preceding  para- 
graph. These  are  simply  the  means  by  the  help  of  which  chemical 
work,  whether  it  be  investigation,  commercial  analysis,  or  manu- 
facturing, may  be  carried  on  more  systematically.  Together  they 
constitute  our  system  for  classifying  the  facts  with  a  view  to  ready 
reference.  The  sample  experiments  (pp.  12-21),  if  reexamined, 
will  show  that  we  there  employed  most  of  the  categories  of  our 
classification  which  have  so  far  been  described. 

Thus,  in  the  experiment  with  iron  and  sulphur  (p.  14),  it  was  first 
our  object  to  find  out  whether  the  bodies  had  interacted  chemically 
on  being  mixed.  To  do  this  we  noted  the  specific  properties  (p.  3*) 
of  the  substances,  separately.  Using  these  properties,  we  were  able 
to  identify  the  same  substances  in  the  mixture,  and  in  this  we  found 
no  substance  with  new  properties.  Part  dissolved  in  carbon  disul- 
phide  and  reappeared  after  evaporation  of  the  solvent  as  yellow, 
rhombic  crystals,  and  the  rest  was  all  magnetic.  In  this  connection 
we  purposely  omitted  all  mention  of  the  quantity  and  temperature, 
because  attributes  (p.  41)  like  the  former  and  conditions  (p.  41)  like 
the  latter  do  not  characterize  substances  (since  they  vary  with  each 
specimen),  and  cannot  be  used  for  identification.  Coincidence  in 
two  or  three  specific  properties  is  generally  sufficient  to  establish 
identity. 

Most  of  the  properties  cannot  be  recognized  readily  in  mixtures, 
as  a  moment's  thought  will  show.  The  general  color  and  the  specific 
gravity  of  a  mixture,  containing  unknown  substances  in  unknown 
proportions,  for  example,  tell  us  little  about  the  corresponding 
properties  of  the  components.  Now,  there  are  few  pure  substances 
in  nature  or  in  the  products  of  experiment,  and  many  mixtures. 
Hence,  separation  of  the  components  of  a  mixture  usually  of  neces- 
sity precedes  the  process  of  identification  just  referred  to.  Thus,  we 
first  removed  the  sulphur  by  dissolving  it  and  then  recovered  it  by 
evaporation  of  the  solvent,  the  boiling-point  of  the  carbon  disulphide 
(46°)  being  much  lower  than  that  of  the  sulphur  (445°)  and  its 
vapor  pressure,  by  virtue  of  which  it  evaporated  at  the  temperature 

*  See  footnote  on  p.  14. 


ENERGY  IN  CHEMICAL  CHANGE 


43 


of  the  room,  being  therefore  relatively  high.  We  secured  the  iron 
because  of  its  insolubility.  Those  specific  properties  which  can  be 
used  for  separating  mixtures,  as  well  as  for  identification,  are  there- 
fore the  most  important  of  all.  They  are  the  melting-point,  solu- 
bility, boiling-point,  specific  gravity,  and  magnetic  qualities,  the  last 
being  applicable  almost  exclusively  to  iron,  however. 


FIG.  20. 

In  connection  with  this  investigation  we  employed  several  of 
the  common  methods  of  manipulation  used  by  the  chemist.  These 
methods  are  derived  from  the  conceptions  described  in  last  paragraph. 
Thus  we  treated  the  mixture  with  a  solvent  (Fig.  9) ,  on  the  assump- 
tion that  if  it  was  heterogeneous  (p.  7)  the  components  would  each 
behave  as  if  alone  present.  We  then  filtered,  a  method  invented  for 
dealing  with  a  heterogeneous  mixture  consisting  of  a  solid  and  a  liquid. 
Decantation  is  often  used  in  such  cases,  when  the  solid  is  specifically 
much  heavier  than  the  solvent  and  settles  readily.  We  allowed  the 
carbon  disulphide  to  evaporate  spontaneously,  and  this  is  our  favor- 
ite method  of  dealing  with  a  mixture  which  is  homogeneous,  and 
therefore  would  run  through  a  filter  as  a  whole  without  suffering 


44  INORGANIC  CHEMISTRY 

separation.  When  the  liquid  has  a  higher  boiling-point  than  50-60° 
C.,  as  water  has,  we  use  heat  from  a  steam-bath  or  Bunsen  flame  to 
promote  the  evaporation.  In  evaporation  we  allow  the  vapor  of  the 
liquid  to  escape,  because  it  is  the  less  volatile,  dissolved  body  that  we 
wish  to  examine.  When  we  desire,  on  the  contrary,  to  examine  the 
liquid,  the  vapor  must  be  condensed.  This  method,  which  we  have 
not  yet  had  occasion  formally  (see  Exercise  5)  to  employ,  is  called 
distillation  (Fig.  20).  The  jacket  round  the  long  tube  is  rilled  with  a 
stream  of  cold  water,  which,  on  account  of  its  high  specific  heat, 
quickly  cools  and  condenses  the  vapor.  The  resulting  liquid  is 
caught  in  a  flask. 

These  methods  may  be  adapted  to  the  investigation  of  any  similar 
problem.  Thus,  gunpowder  is  made  by  the  intimate  mixing  of  sul- 
phur, charcoal,  and  saltpeter  (potassium  nitrate).  If  no  chemical 
interaction  whatever  has  occurred,  a  sample  will  be  wholly  separable 
into  these  components.  If  a  partial  change  has  taken  place,  a  certain 
amount  of  material  with  different  properties  will  be  discovered  in  the 
mixture.  If  the  change  has  been  complete,  no  portion  of  the  original 
substances  will  be  found.  We  must  first  study  the  specific  properties 
of  each  of  the  ingredients  separately,  in  order  that  a  plan  of  sepa- 
ration may  be  devised,  and  that  we  may  have  a  basis  for  comparison 
with  the  products  of  the  separation. 

It  should  be  noted  that  our  methods  of  investigating  the  products 
of  nature  and  of  the  laboratory  are  purposely  limited  so  as  not  to 
separate  chemically  combined,  but  only  physically  mixed  forms  of 
matter.  After  a  physical  individual  has  been  isolated,  and  even  then 
only  if  it  has  new  properties,  and  is  not  at  once  recognized  as  a  known 
substance,  we  next  proceed  to  separate  it  into  its  chemical  constit- 
uents so  as  to  learn  which  constituents  it  contains  and  in  what 
relative  proportions  by  weight. 

The  experiments  with  mercuric  oxide  and  with  silver  nitrate 
simply  rang  the  changes  on  the  same  conceptions,  and  repeated  the 
same  manipulations.  In  every  case  the  specific  properties  —  color, 
crystalline  form,  volatility,  solubility,  and  so  forth  —  by  which  sepa- 
ration and  recognition  were  effected  will  be  found  to  have  been 
mentioned.  When  we  heat  the  mercuric  oxide  (p.  17),*  it  is  resolved 
into  its  constituents.  We  selected  the  experiment  because  the  con- 
ditions are  such  that  separation  of  the  products  from  each  other  and 
from  the  original  material  occurs  spontaneously,  without  further 
manipulation.  The  separation  depends  here  on  the  boiling-points 
*  See  footnote  on  p.  14. 


ENERGY  IN  CHEMICAL  CHANGE  45 

of  the  materials.  That  of  oxygen  is  very  low  (—183°  C.);  so  that, 
being  a  gas  even  at  ordinary  temperatures,  it  comes  off.  That  of 
mercury  is  much  higher  (357°  C.)>  but  the  heat  produced  by  the 
flame  (1000°- 1200°  C.  in  the  bottom  of  the  test-tube)  is  more  than 
sufficient  to  vaporize  this  product  also.  The  speedy  condensation 
on  the  cooler  part  of  the  tube,  causing  separation  of  the  two  products, 
is  due  to  the  wide  difference  in  boiling-points.  The  part  of  the 
mercuric  oxide  which  is  still  unchanged,  on  the  other  hand,  is  entirely 
involatile  even  at  1200°  C.;  so  that  it  does  not  mingle  with  the 
vaporized  products  at  all. 

In  the  action  of  silver  nitrate  on  sodium  chloride  (p.  19),  it  was 
solubility  that  furnished  the  means  of  separating  the  products.  The 
silver  chloride  is  practically  insoluble  in  water,  while  sodium  nitrate 
is  very  soluble.  To  recover  the  latter,  the  boiling-points  were  then 
considered,  and  the  more  easily  vaporized  water  was  driven  off. 

When,  following  the  latter  experiment,  the  silver  chloride  was  ex- 
posed to  light,  resolution  into  silver  and  chlorine  took  place.  The 
products  became  separated  spatially  at  once,  because  chlorine  is  a 
gas  (substance  of  low  b.-p.)  and  silver  is  a  solid  (usually  a  substance 
of  high  b.-p.). 

Fact  and  Law.  —  If  the  preceding  pages  be  reexamined,  it  will 
be  seen  that  a  number  of  facts  are  mentioned  in  them.  We  begin 
the  organization  of  knowledge  according  to  the  scientific  method  by 
trying  to  determine  the  facts.  Thus,  we  find  some  specimens  of 
iron  are  variously  colored,  and  some  are  brittle.  Examination 
shows,  however,  that  the  former  peculiarities  are  due  to  paint,  for 
example,  and  the  latter  to  the  presence  of  carbon  and  other  foreign 
materials  in  the  iron  (cast  iron).  Finally,  we  ascertain  the  facts 
that  iron  itself  is  silver-white  and  tough. 

Facts  are  the  ultimate  units  of  the  structure  of  a  science.  Thus,  a  single  iso- 
lated observation,  no  matter  how  accurately  made,  does  not  furnish  us  with  the 
sort  of  fact  that  can  receive  a  permanent  place  in  our  collection.  It  is  only  after 
much  research  and  thought  that  we  can  ascertain  which  are  the  fixed  elements  in 
the  variety  of  experience  in  any  line,  and  so  determine  what  are  the  facts,  in  the 
sense  in  which  we  have  used  the  term. 

Putting  together  statements  or  laws  like  those  appearing  in 
heavy  type  above  (pp.  11,  17,  18),  is  the  second  step.  We  examine 
facts  of  a  like  kind,  or  pertaining  to  like  phenomena,  to  see  whether 
any  general  statement  can  be  made  that  will  cover  some  feature 


46  INORGANIC  CHEMISTRY 

common  to  the  whole  of  them.  For  example,  after  settling  the 
intrinsic  properties  of  several  substances,  and  then  determining  the 
facts  about  the  way  in  which  these  properties  are  affected  under 
certain  circumstances,  we  decide  that,  when  one  substance  gives  two 
entirely  different  substances,  or  one  of  the  other  varieties  of  change 
(p.  21)  occurs,  the  cases  in  which  this  takes  place  shall  constitute 
a  distinct  class,  and  we  call  them  cases  of  chemical  change.  The 
statement  which  in  a  few  words  or  phrases  sums  up  those  features 
of  all  the  phenomena  of  like  kind  which  are  constant,  is  called  a 
generalization,  or  a  law,  or  a  principle  of  the  science.  A  generaliza- 
tion or  law  in  chemistry,  therefore,  is  a  brief  statement  describing 
some  general  fact  or  constant  mode  of  behavior. 

We  must  remember  that  laws  are  only  true  so  long  as  no  facts  in 
conflict  with  them  are  known.  There  are  no  laws  in  nature.  Nature 
presents  materials  and  phenomena  as  she  pleases.  The  laws  are 
parts  of  science,  which  is  made  by  man  (p.  1),  and  is  a  description 
of  natural  facts  as  man  knows  them.  As  we  shall  see  (p.  54),  at 
least  one  undoubted  exception  to  the  law  of  constant  proportion 
has  recently  (1914)  been  discovered,  and  other  exceptions  to  this 
law  will  undoubtedly  turn  up.  A  law  is  true  only  at  the  date  of  its 
formulation  by  one  chemist  and  its  acceptance  by  other  chemists, 
and  only  so  long  as  no  facts  to  the  contrary  become  known. 

Of  course,  the  same  set  of  facts  may  be  viewed  in  many  ways.  Picking  out 
the  relationships  which  are  most  comprehensive  and,  at  the  same  time,  are  best 
fitted  to  form  part  of  still  broader  generalizations,  or  to  take  their  places  alongside 
of  equally  broad  ones,  requires  the  highest  ability.  The  most  important  laws, 
like  that  describing  the  behavior  of  gases  when  compressed  (Boyle's  law),  are 
usually  connected  with  the  names  of  the  men  by  whom  the  relationships  were  dis- 
covered and  the  generalizations  formulated. 

The  word  "formulate,"  as  applied  to  a  law,  is  preferable  to  the  word,  "dis- 
cover." The  latter  is  ambiguous  and  suggests  that  the  law  existed  before  it  was 
found.  Even  the  relation  which  it  puts  into  words,  did  not,  properly  speaking, 
exist,  because  relations  are  picked  out  of  a  complex  by  the  mind,  and  the  particu- 
lar relation  selected  is  a  property  of  the  mind  and  not  of  the  constituents  of  the 
complex  itself. 

The  reader  must  beware  of  the  misconception  that  a  law  enforces  behavior  in 
accordance  with  its  tenor.  The  mere  statement  that  every  piece  of  matter  attracts 
every  other,  cannot  compel  a  stone  to  fall;  nor  can  the  mode  of  behavior  of  one 
falling  stone  persuade  any  other  to  do  likewise.  The  law  is  simply  a  record 
of  what  is  invariably  observed  to  happen. 

It  is,  therefore,  also  very  misleading  if  we  permit  ourselves  to  say  that  Boyle's 
Jaw  "acts"  so  as  to  "cause"  gases  to  behave  in  a  certain  way,  or  that  the  law 


ENERGY  IN  CHEMICAL  CHANGE  47 

"operates"  to  "produce"  a  certain  behavior,  or  that  other  behavior  is  "impos- 
sible" to  the  gas,  or  that  the  law  of  cohesion  "intervenes"  when  the  gas  is  under 
low  pressure,  and  causes  its  behavior  to  "diverge  from  that  required"  by  Boyle's 
law,  or  to  say  that  a  gas  "disobeys"  Boyle's  law.  In  scientific  discussions,  such 
figures  of  speech  are  alone  permissible  as  throw  light  on  the  subject.  Phrases 
like  the  above,  common  as  their  use  is,  have  been  selected  apparently  with  a  view 
.  to  introducing  a  maximum  of  distortion  and  obscurity;  It  is  the  gas  that  "acts" 
and  gives  rise  to  the  making  of  Boyle's  law,  and  the  latter  is  only  an  epitome  of 
the  way  it  acts.  The  "laws"  of  nature  are  not  unchanging,  although  we  are 
often  told  that  they  are.  It  is  the  behavior  in  nature  that  is  unchanging.  The 
laws  which  attempt  to  describe  this  behavior  are  not  natural  laws,  they  are  man 
made.  They  are  continually  amended,  that  they  may  more  closely  describe 
additional  facts  disclosed  by  research.  Behavior  divergent  from  the  "laws  of 
nature"  is  not  only  possible,  but  is  constantly  being  observed.  Such  divergencies 
furnish,  in  fact,  the  commonest  starting  points  for  fresh  investigation.  Thus 
Boyle's  law  was  amended  by  van  der  Waals  (p.  165),  and  the  law  of  octaves  (q.v.) 
was  greatly  modified  by  Mendelejeff.  It  is  not  the  gas  which  "diverges"  from 
our  statement  or  "disobeys"  our  law,  but  our  statement  which  is  proved  by  the 
behavior  of  the  gas  to  be  inaccurate.  Our  procedure,  in  such  cases,  is  always 
more  logical  than  our  language,  for  we  never  attempt  to  cure  the  gas  of  its  error, 
but  always  the  law  itself  by  suitable  modification  in  its  phraseology. 

Physics  in  Chemistry.  —  It  will  be  seen  that  one  cannot  ac- 
complish anything  in  chemistry  without  acquiring  and  using  some 
knowledge  of  physics.  We  measure  quantities  by  means  of  the 
physical  attributes,  weight  and  volume.  We  produce  chemical 
change  by  arranging  the  physical  conditions,  for  example,  by  mixing, 
heating,  or  using  an  electric  current.  Physical  means  are  the  only 
means  we  possess  for  producing,  stopping,  or  modifying  chemical 
changes.  Again,  we  ascertain  whether  a  chemical  change  has  taken 
place  or  not  by  observing  the  physical  properties  of  the  materials 
before  and  after  the  experiment.  Thus,  we  noted  that  the  red, 
powdery  oxide  of  mercury,  when  heated,  gave  a  liquid  metal  and  a 
gas.  All  the  phenomena  of  chemistry  are  physical.  A  phenomenon 
(Gk.  faivuv,  to  show)  is  literally  something  that  is  seen,  or  more 
generally,  something  that  affects  any  of  the  senses.  Observing 
physical  phenomena  is,  therefore,  our  sole  means  of  studying  chemical 
changes.  Chemical  work  is  therefore  entirely  dependent  upon  the 
skilful  use  of  physical  agencies,  and  the  close  observation  of  physical 
phenomena  for  its  success. 

It  is  only  the  inference,  following  the  experiment  and  the  observa- 
tion, that  is  strictly  chemical.  If  one  substance  gives  two  different 
substances,  or  if  two  substances  give  one  different  substance,  for 


48  INORGANIC  CHEMISTRY 

example,  we  infer  that  a  chemical  change  has  occurred.  We  then 
try  to  recognize  the  substances  by  their  properties  and  name  them. 
Changes  like  that  of  ice  into  water,  or  of  water  into  steam,  and 
vice  versa,  are  not  regarded  as  chemical  changes.  These  are  called 
changes  of  the  state  of  aggregation.  The  solid,  liquid,  and  gaseous 
forms  are  different  states  of  the  same  substance.  The  very  language 
we  use  bears  testimony  to  the  universal  acceptance  of  the  view  that- 
change  of  state  does  not  constitute  a  fundamental  alteration.  Thus, 
we  have  solid  lead  and  molten  lead,  air  (the  gas)  and  liquid  air.  In 
the  case  of  water  alone  has  it  been  found  convenient  to  distinguish  the 
three  forms,  ice,  water,  and  steam,  by  separate  names. 

Attempts  to  Distinguish  between  Physics  and  Chemistry. — 

If  the  reader  has  studied  elementary  chemistry  before,  he  must  beware  of  certain 
misconceptions  which  he  may  have  formed  in  regard  to  the  relation  of  physics  to 
chemistry.  For  several  reasons,  we  have  omitted  most  of  the  usual  material  em- 
ployed in  discussing  the  difference  between  chemistry  and  physics.  To  discuss 
the  difference  between  two  things,  when  the  reader  does  not  yet  know  anything 
about  one  of  them  (chemistry),  and  sometimes  knows  little  or  nothing  about 
the  other,  must  necessarily  be  futile.  Then,  too,  the  alleged  difference  between 
physical  and  chemical  phenomena,  as  phenomena,  does  not  exist.  Heating  a 
platinum  wire  and  allowing  it  to  cool,  heating  mercuric  oxide  and  decomposing 
it,  and  an  animal  walking,  when  regarded  as  phenomena  (appeals  to  the  senses) 
are  all  physical.  The  phenomena  presented  in  chemical  change,  in  the  structure 
of  plants  and  animals,  in  geological  formations,  and  in  pure  physics  are  all  phys- 
ical. After  we  have  assigned  certain  groups  of  such  phenomena  to  other  sciences, 
what  is  left  belongs  to  physics  as  such.  But  physics  still  has  to  be  employed  in 
explaining  how  the  limbs  of  an  animal  work,  how  strata  are  deposited,  and  how 
the  mercury  and  the  oxygen  are  separated,  and  how  each  is  recognized  in  the 
practical  study  of  chemistry.  We  cannot  omit  physics  because  it  is  needed  at 
every  turn.  Chemistry  deals  with  the  interpretation  of  certain  classes  of  phys- 
ical phenomena,  geology  with  other  classes,  and  so  forth. 

Some  of  the  current  distinctions  between  chemical  and  purely  physical 
phenomena  deserve  notice.  Some  physical  changes,  like  the  boiling  of  water 
and  the  condensation  of  the  steam,  are  easily  reversible.  But  no  one  ever  saw  a 
piece  of  broken  glass  restored  to  its  original  condition.  Many  toys  are  broken, 
and  few  mended.  In  fact,  some  physical  changes  are  easily  reversible,  and 
others  are  not.  We  shall  see  later  that  many  chemical  changes  are  also  easily 
reversible,  although  many  are  not.  There  is  no  distinction  along  this  line. 
Again,  in  chemical  change,  all  the  physical  properties  are  altered.  But,  when 
water  is  converted  into  steam  or  ice,  all  the  physical  properties  are  different. 
Change  in  all  physical  properties  occurs  in  physics  also.  Still  again,  the  change 
from  iron  to  rust  is  abrupt  or  sharp,  while  the  changes  in  the  properties  of  iron  as 
it  is  gradually  raised  to  higher  and  higher  temperatures  are  gradual.  But, 


ENERGY  IN  CHEMICAL  CHANGE  49 

when  the  iron  melts,  or  the  water  is  changed  into  steam,  the  change  is  very 
sharp.  On  the  other  hand,  when  water  is  heated  from  0°  to  100°,  chemical 
changes  from  (H2O)3  to  (H2O)2  and  H2O  occur  and,  when  ammonium  chloride 
vapor  is  heated,  more  NH3  +  HC1  is  formed  and  less  NH4C1  remains,  yet  no 
sharp  change  in  properties  is  noticeable.  The  absence  of  abrupt  change  is  due 
to  the  fact  that  the  substances  in  the  water  and  the  gases  in  the  vapor  are  com- 
pletely miscible  with  one  another,  so  that  only  gradual  change  can  be  noticed. 
On  the  other  hand,  rust  and  iron,  or  water  and  steam,  are  not  miscible,  and  so 
two  phases  are  produced.  Hence,  the  change  (physical  or  chemical)  is  sharp 
if  a  new  phase  is  produced,  but  gradual  if  the  products  are  miscible  with  the 
original  substances.  Finally,  it  is  sometimes  said  that  physical  changes  are 
produced  by  physical  or  mechanical  means,  and  chemical  changes  by  chemical 
means.  In  point  of  fact  chemical  changes  are  produced  by  mixing  substances, 
or  by  heating  (producing  more  violent  motion  of  the  molecules),  or  by  using  a 
current  of  electricity  (driving  a  stream  of  electrons  into  the  solution),  or  by 
other  physical  and  mechanical  means,  and  no  other  means  exist.  Practical 
chemistry  consists,  therefore,  in  using  physical  means  to  produce  and  to  control 
chemical  change,  and  to  separate  and  examine  the  results.  Hence,  while  studying 
real  chemistry,  one  is  bound  to  use,  and  to  learn  much  about,  practical  physics. 

Where  Physics  and  Chemistry  Meet.  —  In  the  completeness  of  the 
transformation,  and  in  the  fact  that  only  one  original  substance  and  one  product 
are  required,  the  physical  changes  of  the  nature  of  melting  a  solid  and  vaporizing 
a  liquid  resemble  the  fifth  variety  of  chemical  change.  Where,  then,  is  the  line 
between  chemistry  and  physics  to  be  drawn?  It  is  in  this  fifth  group  only  that  the 
difficulty  is  encountered.  All  agree  that  warm  solid  phosphorus  is  chemically 
identical  with  cold  phosphorus.  Nearly  all  scientific  men  at  present  assign  the 
study  of  the  melting  and  the  vaporization  of  phosphorus  and  other  substances 
to  physics.  Some  chemists  consider  the  solution  of  phosphorus  in  carbon  disul- 
phide  or  some  other  solvent  (it  is  practically  insoluble  in  water),  although  the 
material  is  recovered  unchanged  by  evaporation  of  the  liquid,  as  a  chemical 
change.  But  the  great  majority  regard- this  as  physical  also.  If  not  chemically 
different,  the  solid,  liquid,  gaseous,  and  dissolved  forms  of  a  substance  must  be 
classed  as  mere  physical  states  of  aggregation.  On  the  other  hand,  red  phos- 
phorus is  held  by  most  chemists  to  be  chemically  different  from  yellow  phos- 
phorus. Many  kinds  of  matter  show  as  much  variety  in  form  as  phosphorus, 
and  some  show  more. 

In  solving  an  ordinary  puzzle,  we  work  knowing  that  some  simple  solution 
exists.  When  studying  nature,  we  are  saved  much  embarrassment  by  remember- 
ing that,  in  cases  like  this,  we  are  not  seeking  for  a  clear-cut  distinction  of  whose 
existence  we  are  assured  in  advance.  Nature  is  under  no  obligation  to  furnish 
easily  classifiable  facts  at  all.  The  limitations  of  our  minds  compel  us  to  classify 
as  far  as  we  are  able  to  do  so.  On  this  plan  alone  can  we  master  the  infinity  of 
detail.  But  the  resulting  system  exists  in  our  own  minds,  where  it  originated,  and 
not  in  nature.  In  the  present  instance,  we  are  seeking,  as  always,  to  make  a 


50  INORGANIC  CHEMISTRY 

distinction  for  our  own  convenience.  But,  at  present,  the  effort  to  make  a  final 
distinction  that  can  be  used  consistently  causes  more  inconvenience  than  the  alter- 
native of  letting  the  matter  rest.  The  history  of  previous  experiences  leads  us  to 
hope  that,  with  fuller  knowledge,  we  shall  be  able,  sooner  or  later,  to  construct  a 
system  of  classification  even  for  this  obscure  region.  After  the  distinction  has 
been  made  and  accepted  by  chemists,  and  not  till  then,  will  it  be  possible  to  explain 
it  to  students. 

Summary.  —  In  this  chapter  we  have  added  considerably  to 
our  conception  of  the  scope  of  chemistry  (cf.  pp.  13,  26) .  Although 
our  survey  is  by  no  means  yet  complete,  we  may  condense  our  results 
as  follows: 

Chemistry  deals,  not  only  with  the  changes  in  composition  and 
constitution  which  substances  undergo,  but  also  with  the  transforma- 
tions of  energy  which  accompany  them.  To  convert  the  isolated  facts 
into  a  science  we  classify  related  parts  under  laws,  such  as  that  of 
conservation  of  energy  (p.  32),  and  under  conceptions,  such  as  those 
of  internal  energy  (p.  34),  and  chemical  activity  (p.  37).  We  also 
distinguish  between  specific  physical  properties,  attributes  and  con- 
ditions (pp.  40-41).  In  the  following  paragraphs  we  have  indicated 
briefly  the  use  to  which  these  conceptions  and  this  classification  are 
put.  Finally,  we  have  discussed  the  meaning  of  terms  like  cause, 
fact,  and  law,  and  the  relation  of  physics  to  chemistry. 

Chemical  laboratory  work  consists  largely  in  the  separation,  rec- 
ognition, and  description  of  substances.  The  importance,  especially, 
of  thorough  familiarity  with  specific  properties  and  the  influence  of 
conditions  (for  example,  temperature)  to  these  ends  is  shown  by  the 
examples  (pp.  42^45).  The  system  of  classification  as  a  whole  is  part 
of  the  everyday  mode  of  thought  of  the  chemist,  for  thought  consists 
largely  in  comparing  and  contrasting,  and  our  system  of  classification 
furnishes  the  plan  of  this  so  far  as  chemistry  is  concerned.  Learning 
chemistry  consists,  therefore,  in  large  part  in  learning  this  classifica- 
tion and  becoming  habituated  to  its  use. 

The  influence  of  conditions  has  as  yet  been  barely  touched.  It 
will  be  dealt  with  more  explicitly  as  occasion  offers.  The  attribute  of 
quantity,  which  has  already  received  some  attention  (p.  13),  will  form 
the  basis  of  discussion  in  the  next  chapter. 

Exercises.  —  1.  What  is  the  original  form  of  energy  used  in, 
(a)  a  hydro-electric  plant,  (b)  an  automobile,  (c)  a  watch,  (d)  a  clock 
with  weights? 


ENERGY  IN  CHEMICAL  CHANGE  51 

2.  From  what  form  of  energy  comes,  (a)  the  heat  of  a  burning 
candle,  (b)  the  heat  in  an  electric  light,  (c)  the  electricity  of  a  battery? 

3.  Describe,  (a)  a  red-hot  rod  of  iron,  10  cm.  long  by  1  cm. 
diameter,  weighing  58.5  g.,  (b)  a  solution  of  5  g.  of  sulphur  in  20  c.c. 
of  carbon  disulphide  at  18°  C.     In  doing  so,  divide  the  description 
into  attributes,  conditions,  and  properties. 

4.  Color,  volume,  melting-point,  density,  temperature,  weight, 
crystalline  form,  boiling-point,  pressure.     In  respect  to  which  of 
these  qualities  will,  (a)  one  substance  differ  from  another,  (b)  one 
specimen  of  the  same  substance  differ  from  another,  (c)  one  specimen 
of.  the  same  substance  differ  at  different  times?     Give  the  technical 
names  and  uses  of  the  sets  of  qualities  which  fall  under  a,  6,  and  c, 
respectively. 

5.  What  properties  are  essential  in  order  that  two  substances 
may  be  separated  by,  (a)  evaporation,  (6)  nitration,  (c)  distillation? 

6.  What  is  the  cause,  (a)  of  the  precipitation  of  silver  chloride 
(p.  19),  (b)  of  the  union  of  iron  and  sulphur  (p.  16)? 


CHAPTER  III 
COMBINING  PROPORTIONS  BY  WEIGHT 

IF  we  now  return  to  the  illustrations  of  chemical  phenomena 
which  we  have  been  studying  (pp.  14-21),  we  shall  find  a  new  question 
arising  naturally  out  of  them.  This  is,  whether  the  mass  of  the  ma- 
terials is  altered,  as  are  the  other  attributes,  in  these  chemical  changes. 

Conservation  of  Mass:  Fourth  Characteristic  of  Chemical 
Change.  —  The  most  painstaking  chemical  work  seems  to  show 
that,  if  all  the  substances  concerned  in  the  chemical  change  are 
weighed  before  and  after  the  change,  there  is  no  evidence  of  any 
alteration  in  the  quantity  of  matter.  The  two  weights,  representing 
the  sums  of  the  constituents  and  of  the  products  respectively,  are, 
indeed,  never  absolutely  identical,  but  the  more  careful  the  work  and 
the  more  delicate  the  instrument  used  in  weighing,  the  more  nearly 
do  the  values  approach  identity.  We  are  able  to  state,  therefore,  as 
a  law  of  all  chemical  phenomena,  that:  The  mass  of  a  system  is  not 
affected  by  any  chemical  change  within  the  system. 

This  statement  simply  means  that  th  e  great  law  of  the  conserva- 
tion of  mass  holds  true  in  chemistry  as  it  does  in  physics.  Chemical 
changes,  thoroughgoing  as  they  are  in  res  pect  to  all  other  qualities, 
do  not  affect  the  mass;  an  element  carries  with  it  its  weight,  entirely 
unchanged,  through  the  most  complicated  chemical  transformations. 

A  law,  as  we  have  seen  (p.  45),  is  a  condensed  statement  describing  some  con- 
stant mode  of  behavior.  It  is  simply  a  summary  of  our  experience.  As  such,  it  is 
subject  to  modification  when  a  fact  is  discovered  with  which  it  conflicts.  Thus, 
it  is  perfectly  possible  that  we  may  yet  find  cases  of  demonstrable  changes  in 
weight  accompanying  other  physical  or  chemical  changes  in  a  limited  system. 
Indeed,  it  has  more  than  once  been  alleged  that  such  changes  have  been  observed. 
It  used  to  be  a  law  that  the  earth  was  flat.  It  is  now  more  correct  to  say  that  a 
limited  area  of  perfectly  level  ground  is  very  nearly  flat. 

It  will  be  observed  that  the  phrasing  of  the  above  law  carefully  limits  its 
scope  to  amounts  of  matter  such  as  are  dealt  with  in  laboratory  experience.  We 
have  no  evidence  on  which  to  make  any  statements  about  the  mass  of  matter  in 
more  extensive  chemical  changes.  A  common  form  of  the  law,  to  the  effect  that 
"the  mass  of  matter  in  the  universe  is  unchangeable  in  amount,"  is  not  a  law  at 

52 


COMBINING  PROPORTIONS  BY  WEIGHT 


53 


all,  in  the  only  sense  in  which  the  word  is  used  in  science.  It  is  a  statement  in 
regard  to  supposed  facts  which  are  almost  entirely  beyond  our  experience.  It  is, 
therefore,  a  proposition  of  a  transcendental  (that  is,  transcending  experience) 
nature,  and  has  its  proper  place  in  metaphysics.  Astronomical  observation,  it  is 
true,  has  as  yet  furnished  no  evidence  of  changes  in  the  mass  of  our  own  or  other 
celestial  systems.  But,  absence  of  evidence  to  the  contrary,  especially  considering 
the  relatively  limited  scope  of  our  knowledge,  both  in  respect  to  space  and  time,  is 
far  from  being  proof  of  the  correctness  of  the  proposition. 

Superficial  observation,  as  of  a  growing  tree,  might  seem  to  give 
evidence  of  the  very  opposite  of  conservation  of  matter.  But  here 
the  carbon  dioxide  gas  in  the  air,  the  most  important  source  of 
nourishment  for  plants,  is  overlooked.  Similarly  the  gradual  dis- 
appearance of  a  candle  by  combustion  seems  to  illustrate  the  destruc- 
tion of  matter.  But  if  we  insert  sticks  of  sodium  hydroxide  in  a 
U-tube  (Fig.  21)  to  catch  the  gases  which  rise  through  the  flame,  we 
find  that  the  gases  weigh  even  more  than  the  part  of  the  candle  which 


Fia.  21. 


Fio.  22. 


has  been  sacrificed  in  making  them.  When  we  take  account  of  the 
weight  of  the  oxygen  obtained  from  the  air  which  sustains  the  com- 
bustion, we  find  that  there  is  really  neither  loss  nor  gain  in  weight. 
If  we  carry  out  chemical  changes  in  closed  vessels  (Fig.  22),  which 
permit  neither  escape  nor  access  of  material,  we  find  that  the  weight 
does  not  alter. 

One  way  of  stating  the  difference  between  chemistry  and  physics  is  to  say  that 
changes  in  which  both  the  mass  and  the  identity  of  the  substance  are  conserved 
belong  to  the  latter,  while  those  in  which  the  mass  alone  is  conserved  belong  to 
chemistry. 


54  INORGANIC  CHEMISTRY 

The  Law  of  Definite  Proportions:  Fifth  Characteristic  of 
Chemical  Change.  —  In  making  a  chemical  compound,  may  we 
use  varying  proportions  of  the  constituents?  The  controversy  be- 
tween Berthollet  and  Proust,  in  which  the  former  supported  the  affirm- 
ative and  the  latter  the  negative  side  of  this  question,  was  one  of  the 
chief  features  of  the  chemical  history  of  the  early  part  of  the  nine- 
teenth century.  The  ways  of  forming  or  decomposing  a  compound, 
or  of  carrying  out  a  more  complex  chemical  change,  may  be  varied 
indefinitely.  The  apparatus,  the  mode  of  experiment,  and  the  pro- 
portions of  the  materials,  may  be  altered  at  our  will.  But,  up  to 
1914,  in  spite  of  an  enormous  amount  of  careful  work,  no  case  of 
variation  in  the  proportion  of  the  constituents  actually  used  or  pro- 
duced in  a  given  chemical  action  has  come  to  light.  If  too  much 
of  one  constituent,  for  example,  is  taken,  a  part  simply  remains 
unchanged.  A  higher  temperature  may  hasten  the  chemical  action, 
but  it  does  not  affect  the  quantitative  composition  of  the  products, 
provided  the  resulting  substances  are  of  the  same  nature.  It  was 
the  work  of  Stas  (1860-65)  which  settled  the  question,  by  proving 
that  even  slight  variations  cannot  be  detected,  and  disposed  of 
Berthollet 's  objection's.  It  is,  therefore,  a  characteristic  of  chemical 
phenomena  that :  In  every  sample  of  each  compound  substance  formed 
or  decomposed,  the  proportion  by  weight  of  the  constituents  is  always 
the  same.  This  statement  of  i'act  is  known  as  the  law  of  definite,  or 
of  constant  proportions.  When  the  composition  of  a  substance  seems 
to  be  variable,  it  is  usually  found  on  closer  examination  that  mechani- 
cal mixtures  of  some  kind  were  being  mistaken  for  pure  substances. 

Another  form  of  statement,  which  is  a  corollary  of  this  one,  and  is 
applicable  more  directly  to  complex  chemical  actions,  is:  The  ratio 
by  weight  of  any  one  of  the  factors  or  products  of  a  chemical  change 
to  any  other  is  constant. 

One  exception  to  this  law  has  been  discovered.  Ordinary  lead 
chloride  contains  lead  and  chlorine  in  the  proportion  207.2  :  2  X 
35.46.  Richards,  of  Harvard,  however,  has  found  that  the  lead 
contained  in  uranium  ores  gives  a  chloride  in  which  the  proportion 
of  lead  is  from  206.1  to  206.8  to  the  same  weight  of  chlorine.  Honig- 
schmidt  and  Horowitz  found  lead  from  the  same  general  source  to 
give  in  some  cases  proportions  as  low  as  206.05  and  206.06,  yet  this 
lead  gave  the  same  spectrum  as  ordinary  lead.  Richards  found  this 
lead  to  have  a  density  slightly  lower  (11.288)  than  that  of  ordinary  lead 
(11.337).  Again,  Soddy  found  that  lead  extracted  from  thorium 
ores  gave  a  chloride  containing  a  proportion  of  lead  higher  than  the 


COMBINING  PROPORTIONS  BY  WEIGHT  55 

ordinary  ore,  namely,  208.4  :  2  X  35.46.  There  thus  appear  to  be 
three  chlorides  of  lead,  having,  so  far  as  we  know,  the  same  properties, 
and  being  therefore  the  same  substance,  yet  differing  in  composition. 
This  illustrates  the  fact  that  a  law  is  true  only  until  facts  at  variance 
with  it  are  encountered.  This  law  was  absolutely  true  up  to  1914, 
but  is  now  true  of  all  compounds  with  the  exception  of  those  of  lead 
extracted  from  the  ores  of  uranium  and  thorium.  Still  other  excep- 
tions will  undoubtedly  be  discovered  very  soon. 

The  Measurement  of  Combining  Proportions.  —  The  most 
exact  measurement  of  the  proportions  in  which  the  elements  com- 
bine to  form  compounds  involves  manipulations  too  elaborate  to 
be  gone  into  here.  Operations  of  the  same  nature  are  described 
in  works  on  quantitative  analysis.  One  or  two  brief  statements, 
diagrammatic  rather  than  accurate,  will  show  the  principles, 
however. 

If  we  take  a  weighed  quantity  of  iron  in  a  test-tube  and  heat  it  with 
more  than  enough  sulphur  (an  excess  of  sulphur) ,  we  get  free  sulphur 
along  with  the  ferrous  sulphide  (p.  16),  and  no  free  iron  survives. 
We  may  remove  the  free  sulphur  by  washing  the  solid  with  carbon 
disulphide.  The  difference 

between  the  weights  of  fer-    C=CHZ  — k  ^wfa ' ' — 

rous  sulphide  and  iron  gives 
the  amount  of  sulphur  com- 
bined with  the  known  quan- 
tity of  the  latter. 

As  an  example  of  the 
study  of  rusting,  we  may 
weigh  a  small  amount  of 

copper  in  the  form  of  powder  FJG  23 

in  a  porcelain  boat  and  pass 

oxygen  over  the  heated  metal  (Fig.  23).  The  formation  of  cupric 
oxide  takes  place  rapidly.  If  we  limit  the  oxygen,  part  of  the  copper 
may  remain  unaltered;  if  we  use  it  freely,  the  excess  will  pass  on 
unchanged.  A  given  weight  of  copper  cannot  be  induced  to  take  up 
more  than  a  certain  amount  of  oxygen,  and  use  of  a  less  amount 
simply  limits  the  amount  of  copper  transformed  into  oxide.  The 
original  weight  of  the  copper,  and  the  increase  in  weight,  representing 
oxygen,  give  us  the  data  for  determining  the  composition  of  cupric 
oxide.  The  data  furnished  by  one  rough  lecture-experiment,  for 
example,  were  as  follows: 


56 


INORGANIC   CHEMISTRY 


Weight  of  boat  +  copper 4.278  g. 

Weight  of  boat  empty 3.428  g. 

Difference  =  weight  of  copper 0.850  g. 

Weight  after  addition  of  oxygen 4.488  g. 

Weight  without  oxygen 4.278  g. 

Difference  =  weight  of  oxygen 0.210  g. 

The  proportion  of  copper  to  oxygen,  so  far  as  this  one  measurement 
goes,  is  therefore  85  :  21. 

The  results  of  quantitative  experiments  are  usually  recorded  in 
the  form  of  parts  in  one  hundred.  To  find  the  percentage  of  each 
constituent,  we  observe  that  the  proportion  of  copper  is  85:  85+21, 
or  TW  of  the  whole.  That  of  the  oxygen  is  TV#  of  the  whole.  Thus 
the  percentages  are: 

Copper,         106  :  85  ::  100  :  x          x  =  80.2 
Oxygen,         106  :  21  ::  100  :  x'        x'  =  19.8 

Naturally,  the  mean  of  the  results  of  a  number  of  more  carefully 
managed  experiments  will  be  nearer  the  true  proportion.  The  per- 
centages at  present  accepted, 
as  most  accurate  are  79.9  and 
20.1. 

In  the  case  of  mercuric 
oxide,  we  may  decompose  a 
known  weight  of  the  oxide  (p. 
17)  and  afterwards  weigh  the 
mercury  and  ascertain  the  oxy- 
gen by  difference. 

Finally,  a  strip  of  the  metal 
magnesium  may  be  set  on  fire 
in  the  air.  It  gives  out  a  daz- 
zling white  light  in  burning, 
and  on  this  account  the  pow- 
dered metal  is  used  in  making 
flash-light  powder  for  photog- 
raphy. The  product  is  mag- 
nesium oxide,  a  white  sub- 
stance, which  partly  rises  as  a  dense  smoke  and  partly  falls  on  the 
ground.  In  a  loosely  closed  porcelain  vessel  (Fig.  24)  the  metal  may 
be  burned  slowly,  with  the  help  of  the  heat  from  a  small  flame,  and 
the  oxide  may  be  retained.  The  weight  of  magnesium  ribbon  taken 


Fia.  24. 


COMBINING  PROPORTIONS  BY  WEIGHT  57 

and  the  increase  in  weight  due  to  oxygen  give  the  data  for  calculating 
the  proportions  of  the  constituents. 

The  following  figures  show  the  results  of  some  experiments  like 
these,  and  represent  the  percentage  composition  of  the  products. 
Only  two  places  of  decimals  are  given  in  each  case,  the  numbers 
following  the  second  decimal  being  omitted,  and  hence  the  total,  in 
one  instance,  appears  to  be  only  99.99.  The  numbers  in  parenthesis 
will  be  explained  presently: 

(1)   Cupric  oxide  (2)    Mercuric  oxide 

Copper          79.9        ["31.81  Mercury      92.61         ri00.3"| 

Oxygen          20.1        L  8    J  Ogygen          7.39        L    8     J 

(3)   Water  (4)    Chlorine  monoxide 

Hydrogen     11.19      p. 0081  Chlorine      81.8          r35.46n 

Oxygen         88.81      L    8    J  Oxygen       18.2  L    8     J 

Combining  Weights.  —  The  percentages  in  the  above  list  rep- 
resent the  true  proportions  by  weight  in  the  various  compounds,  but 
naturally  the  individual  numbers  constituting  those  proportions  have 
no  chemical  significance  whatever.  They  are  arbitrary  values 
selected  so  that  the  two  factors  in  the  proportion  may  together  make 
100.  Each  pair  represents  the  constant  ratio  which  is  the  mean 
result  of  numerous  experiments. 

We  begin  the  effort  to  reduce  these  numbers  to  order  by  selecting 
one  element  as  our  starting-point,  and  by  taking  some  convenient 
weight  of  it  as  the  basis.  As  we  shall  see,  it  makes  no  difference  what 
choice  we  make  in  either  respect.  To  avoid  waste  of  time,  we  shall, 
therefore,  use  oxygen,  as  it  is  the  element  generally  preferred  by 
chemists  for  the  purpose.  '  The  reason  for  this  preference  will  be 
apparent  later  (see  pp.  61,  66). 

We  should  naturally  be  inclined  to  use  1  part  of  the  element  as  our 
basis.  But  our  later  steps  involve  finding  out  what  amounts  of  the 
other  elements  combine  with  this  quantity,  and  we  perceive  that  the 
amount  in  the  case  of  hydrogen  will  be  only  0.126  parts.  We  calcu- 
late this  from  (3):  88.81  :  11.18  ::  1  :  x  (=  0.126).  If,  however,  we 
take  8  parts  of  oxygen,  this  amount  of  hydrogen  is  also  increased  eight 
times  and  becomes  1.008.  As  no  element  is  found  to  combine  in 
smaller  proportions  than  hydrogen,  we  are  satisfied  that  a  scale  for  our 
numbers  based  on  8  parts  of  oxygen  will  not  involve  any  values  less 
than  1.  The  choice  of  scale  is  purely  one  of  convenience. 

We  now  proceed  to  calculate  from  the  data  given  auove  the  weight 
of  each  of  the  other  four  elements  which  combines  \  ith  8  parts  of 


58  INORGANIC  CHEMISTRY 

oxygen.  From  (1)  we  calculate  this  weight  of  copper,  thus,  20.1: 
79.9  ::  8  :  x  (=  31.8  parts  of  copper).  Similarly  we  find  that  8 
parts  of  oxygen  combine,  in  (2)  with  100.3  parts  of  mercury,  in  (3) 
with  1.008  parts  of  hydrogen,  and  in  (4)  with  35.46  parts  of  chlorine. 
Oxygen  unites  with  almost  all  the  known  elements,  and  these  four 
compounds  have  been  chosen  simply  as  a  sample  group. 

OXYGEN,  8 


COPPER,  31.8    MERCURY,  100.3    HYDROGEN,  1.008    CHLORINE,  35.46 

Now  chlorine  combines,  not  only  with  oxygen,  but  also  with 
copper,  mercury,  and  hydrogen,  and  measurement  shows  that  the 
combining  proportions  are  represented,  exactly,  by  the  very  same  numbers 
as  before.  From  the  two  independent  facts  that  8  parts  of  oxygen 
combine  with  31.8  parts  of  copper  and  with  35.46  parts  of  chlorine, 
we  could  not  possibly  have  foretold  the  proportions  in  which  copper 
and  chlorine  would  combine  with  one  another.  Yet  measurement 
shows  it  to  be  31.8  :  35.46  precisely.  In  the  following  table  the 
proportions  in  which  the  elements  combine  with  chlorine  are  placed 
under  the  corresponding  parts  of  the  names  of  the  compounds  with 
chlorine : 

CUPRIC  CHLORIDE  MERCURIC  CHLORIDE  HYDROGEN  CHLORIDE 

31.8  :  35.46  100.3  :  35.46  1.008  :  35.46 

If  additional  elements  had  been  included  in  the  group,  a  combining 
number  could  have  been  found  for  each,  and  seeming  coincidences  of 
the  same  nature  would  have  multiplied  rapidly.  Thus,  sulphur 
unites  with  hydrogen  to  give  hydrogen  sulphide.  If,  to  maintain  the 
same  scale,  we  use  1.008  parts  of  hydrogen  in  expressing  the  propor- 
tion, we  find  that  the  combining  ratio  is  1.008  :  16.03.  This  result 
could  not  enable  us  to  predict  the  proportion  of  copper  to  sulphur  in 
cupric  sulphide,  but  measurement  shows  it  to  be  31.8  :  16.03.  Again, 
mercury  and  sulphur  unite  in  the  proportion  100.3  :  16.03. 

SULPHUR,  with  the  value  16.03,  may  therefore  be  added  to  the 
series  of  equivalent  weights. 

Law  of  Multiple  Proportions.  —  One  other  remarkable  fact 
remains  to  be  noted.  There  are  two  different  compounds  of  copper 
with  oxygen.  Cuprous  oxide  (q.v.),  the  one  not  mentioned  hitherto, 


COMBINING  PROPORTIONS  BY  WEIGHT 


59 


is  found  in  nature  as  a  dark-red  mineral  which  is  entirely  different 
from  cupric  oxide  in  physical  properties.  It  can  also  be  prepared 
in  the  laboratory,  but  not  by  simply  passing  oxygen  over  heated 
copper.  Now,  analysis  shows  that  in  cuprous  oxide  the  proportion 
of  copper  combined  with  8  parts  of  oxygen  is  63.6.  This  new  number 
for  copper  is  not  unrelated  to  the  corresponding  value  in  cupric 
oxide  (triz.j  31.8),  for  it  is  exactly  twice  as  great.  Again  mercury  forms 
mercurous  oxide  (see  below)  as  well  as  mercuric  oxide,  and  in  the 
former  the  proportion  of  oxygen  to  mercury  is  8  :  200.6.  The 
proportions  of  mercury  uniting  with  8  parts  of  oxygen  are  there- 
fore 100.3  and  200.6,  and  no  other  proportions  are  known.  Still  again, 
1.008  parts  of  hydrogen  unite  with  8  parts  of  oxygen  in  water  and 
with  2X8  parts  of  oxygen  in  hydrogen  peroxide.  The  fact  sug- 
gested by  these  three  examples  is  a  general  one.  It  was  discovered 
by  Dal  ton  (1804)  and  was  embodied  by  him  in  a  statement  known 
as  the  law  of  multiple  proportions,  which  ran  somewhat  as  follows: 
If  two  elements  unite  in  more  than  one  proportion,  forming  two  or 
more  compounds,  the  quantities  of  one  of  the  elements,  which  in  the 
different  compounds  are  united  with  identical  amounts  of  the  other, 
stand  to  one  another  in  the  ratio  of  integral  numbers,  which  are  usually 
small. 

The  Law  of  Combining  Weights:  Sixth  Characteristic. — 

The  reader  should  now  examine  carefully  the  following  table  of 
combining  proportions.  It  includes  all  the  compounds  made  up  of 
pairs  of  the  six  sample  elements  under  consideration,  so  far  as  the 
existence  and  composition  of  such  compounds  have  been  deter- 
mined with  certainty.  The  substances  in  black  type  are  the  ones 
from  whose  composition  we  originally  derived  the  combining  numbers, 
the  others  illustrate  the  uniform  recurrence  of  the  same  numbers: 


Cupric  oxide 

31.8  :  8 
Cuprous  oxide 

2X31.8  :8 
Cupric  sulphide 

31.8  :  16.03 
Cuprous  sulphide 

2X31.8  :  16.03 
Cupric  chloride 

31.8  :  35.46 


Mercuric  oxide 

100.3  :  8 
Mercurous  oxide 

2X100.3  :8 
Mercuric  sulphide 

100.3  :  16.03 
Mercurous  sulphide 

2X100.3  :  16.03 
Mercuric  chloride 

100.3  :  35.46 


Hydrogen  monoxide  (water) 

1.008  :  8 
Hydrogen  peroxide 

1.008  :2X8 
Hydrogen  sulphide 

1.008  :  16.03 


Hydrogen  chloride 
1.008  : 35.46 


60 


INORGANIC  CHEMISTRY 


Cuprous  chloride 
2X31.8  :  35.46 

Sulphur  dioxide 
16.03  :  2X8 

Sulphur  trioxide 
16.03  :3X8 


Mercurous  chloride 
2X100.3  :  35.46 

Chlorine  monoxide 
35.46  :  8 

Chlorine  dioxide 
35.46  :  4X8 


Sulphur  monochloride 
2X16.03  :  35.46 

Sulphur  tetrachloride 
16.03  :  2X35.46 


It  will  be  observed  that  the  same  number  serves  always  to  express 
the  combining  proportions  of  a  given  element,  provided  that  multi- 
plication by  an  integer  is  permitted  when  necessary.  There  are  also 
some  combinations  of  three  of  the  same  elements  in  one  compound. 
But  even  in  such  cases  the  fundamental  numbers  still  reappear. 
Thus  oxygen,  chlorine,  and  hydrogen  combine  in  the  proportions: 

and  8X8:35.46:1.008. 


2X8:35.46:  1.008, 


6X8:35.46:1.008, 


Nor  is  the  recurrence  of  a  fundamental  number  an  exclusive  property 
of  the  six  elements  we  have  chosen  for  illustration.  A  vast  table 
containing  every  known  element,  and  every  compound  of  every  ele- 
ment, if  prepared  in  the  same  way  as  that  given  above  (by  starting 
with  a  fixed  number  for  one  element  and  calculating  the  combining 
proportions  in  all  the  compounds  with  this  number  as  a  basis) ,  would 
show  that  every  element  uses  an  individual,  fundamental  number, 
or  integral  multiples  of  this  number,  in  every  one  of  its  combinations. 
No  exceptions  to  this  rule  would  be  found  in  the  whole  mass  of  data. 
These  relations  become  clearer  when  represented  diagrammatically. 
Thus  the  five  elements,  omitting  oxygen,  give  the  following  com- 
pounds and  no  others : 


1:1 


1:2- 


-2:1 


Copper 1:1 — Sulphur — 1:1 — Mercury — 1:1 — Chlorine — 1:1 — Hydrogen 

(31.8)  (16.03)  (100.3)  (35.46)  (1.008) 


.1 

.1 

1  .1              

It  will  be  observed  that  hydrogen  forms  stable  compounds  with  only  two  of 
the  other  four  elements  in  the  series.     If,  however,  oxygen  had  been  included, 


COMBINING   PROPORTIONS  BY  WEIGHT  61 

compounds  of  oxygen  with  all  the  other  five  would  have  demanded  recognition. 
This  illustrates  the  reason  given  below  for  the  preference  of  oxygen  as  the  funda- 
mental element. 

The  law  of  combining  weights  may  be  put  briefly  thus:  In  every 
compound  substance,  the  proportion  by  weight  of  each  element  may 
be  expressed  by  a  fixed  number,  one  for  each  element,  or  by  a  multi- 
ple of  the  number  by  some  integer  (whole  number).  It  describes 
what  is  perhaps  the  most  striking  of  all  the  characteristics  of  chemi- 
cal action.  Clearly  it  does  not  apply  to  mixtures,  for  any  irregular 
proportion  could  be  used  in  the  physical  process  of  mixing. 

It  is,  perhaps,  hardly  necessary  to  point  out  that  the  law  of  multiple  pro- 
portions is  simply  a  partial  statement  whose  whole  content  is  included  in  this 
far  greater  generalization.  No  one  chemist  succeeded  in  discovering  this  property 
of  combining  weights.  The  work  of  J.  B.  Richter,  Dalton,  and  many  others 
contributed  to  it. 

Most  of  the  first  exact  determinations  of  combining  proportions  were  made  by 
Berzelius  before  1830.  It  should  be  added  that,  while  the  combining  weights,  with 
the  exception  of  that  of  oxygen  which  is  the  standard,  are  never  actually  whole 
numbers  (although  they  often  approach  such  integral  values),  the  integers  used  to 
multiply  them,  when  they  are  employed  to  express  combining  proportions,  are  so 
most  exactly.  Even  in  determinations  by  methods  of  the  highest  refinement,  the 
factors  to  be  used  in  multiplying  the  combining  weights  are  always  found  to 
diverge  from  whole  numbers  by  amounts  within  the  known  errors  of  the  method 
of  measurement. 

Reexamination  of  Combining  Weights.  —  If  the  reader  will 
now  reexamine  carefully  the  way  in  which  the  data  were  handled,  the 
following  significant  facts  will  be  noted: 

1.  Oxygen  was  made  the  starting-point  of  the  system  and  the 
value  8  was  assigned  to  its  combining  weight.*  Had  a  different  value 
been  used,  all  the  numbers  would  have  been  different.  But  the 
change  would  have  affected  all  the  numbers  in  the  same  proportion 

*  Oxygen  was  chosen  as  the  basis  of  the  system  because  the  exact  deter- 
minations of  the  combining  weights  of  most  of  the  elements  have  actually  been 
made  by  direct  union  with  oxygen  or  with  the  help  of  but  one  intermediate 
step.  If  the  question  had  been  one  of  mathematics,  hydrogen,  the  element 
with  the  lowest  combining  proportions,  would  have  furnished  the  basis  and 
unit  of  the  scale.  But  the  question  was  the  practical  one  of  getting  the  most 
accurate  measurements  for  the  relative  magnitudes  of  the  numbers,  so  oxygen 
was  chosen  instead.  Nevertheless,  the  value  8  was  selected  in  order  that  the 
advantage  of  having  a  mathematical  unit,  or  something  close  to  it,  in  the  com- 
bining weight  of  hydrogen,  might  be  retained  also. 


62  INORGANIC  CHEMISTRY 

and  so  only  the  scale  of  the  numbers  would  have  been  changed.  An 
individual  combining  number,  one  for  each  element,  would  have 
still  recurred  wherever  the  element  itself  appeared. 

2.  Even  use  of  another  element  as  the  initial  one  would  not  have 
prevented  the  discovery  of  the  law.     Thus  with  hydrogen  as  the 
initial  element,  and  the  value  1.008,  no  change  whatever  would  have 
been  noted.     With  a  different  value  for  hydrogen,  a  change  of  scale 
in  all  the  numbers  would  have  followed,  but  individual  combining 
weights  would  have  appeared  as  before. 

3.  Finally,  if  cuprous  oxide  had  been  used  instead  of  cupric  oxide 
for  the  first  proportion,  the  value  found  for  copper  would  have  changed 
to  63.6,  while  the  other  numbers  would  have  remained  unaffected. 
But  this  number  would  serve  as  the  combining  weight  of  copper  just 
as  well  as  31.8,  for  the  composition  of  cupric  oxide  can  be  expressed 
by  the  proportion  63.6  :  2X8  as  well  as  by  31.8  :  8,  and  that  of 
cupric  sulphide  by  63.6  :  2X16.03  as  well  as  by  31.8  :  16.03.     An 
important  conclusion  therefore  follows  from  these  considerations,  and 
we  shall  have  occasion  to  use  it  presently,  namely,  that  any  one  or 
more  of  the  equivalent  weights  (with  scale  oxygen  =  8)  may  be 
separately  multiplied  by  an  integer  without  its  usefulness  as  a  member 
of  the  series  being  at  all  impaired. 

The  importance  of  the  fact  described  in  the  law  of  combining 
weights  cannot  be  emphasized  too  strongly.  Without  this  fact,  the 
remembering  of  the  compositions  of  chemical  substances,  necessary 
as  it  is  to  the  chemist,  would  have  been  completely  beyond  the  power 
of  any  ordinary  memory.  With  it,  the  task  becomes  comparatively 
simple.  It  is  only  necessary  to  decide  on  the  best  system  of  values 
for  the  combining  weights,  and  then,  regarding  the  value  of  this  for 
each  element  as  the  unit  of  weight  for  that  element,  to  express  the  pro- 
portions of  the  element  in  every  compound  by  the  proper  multiples. 
Thus,  given  a  list  of  the  combining  weights,  one  for  each  element,  only 
the  small  integral  multiples  have  to  be  kept  in  mind  in  connection 
with  each  compound. 

The  reader  will  require  a  little  time,  however,  before  he  becomes 
accustomed  to  the  use,  not  of  a  single  unit  of  weight,  but  of  a  different 
one  for  each  element.  Chemistry  is  the  only  science  in  which  the 
physical  unit  of  weight,  which  is  the  same  for  all  materials,  is  not 
employed  for  every  purpose.  The  physical  manipulations  of  the 
chemist  are  carried  out  with  the  use  of  physical  units,  but  the  chemi- 
cal results  are  expressed  in  terms  of  individual  unit  quantities  of  the 
several  elements,  the  combining  weights. 


COMBINING  PROPORTIONS  BY  WEIGHT 


63 


The  individual  units  actually  used  for  each  element  are  not  in  all 
cases  identical  with  those  we  have  given.  The  final  values  will  be 
discussed  in  the  next  two  sections. 

Equivalent  Weights.  —  The  phrase  "combining  weights"  is  a 
general  one,  referring  to  any  values  that  can  be  employed  in  express- 
ing the  proportions  by  weight  used  in  combining,  thus,  for  oxygen, 
the  values  8,  16,  24,  or  32,  or  any  other  multiple  of  8  would  serve  the 
purpose.  Not  only  so,  but  we  could  arbitrarily  assign  any  number 
we  chose  to  oxygen,  such  as  3  or  3.14159,  and  call  it  the  combining 
weight,  provided  we  changed  the  quantities  of  other  elements  uniting 
with  oxygen  correspondingly.  In  other  words,  combining  weights 
have  no  established  values  in  chemistry.  The  term  is,  therefore, 
suited  for  use  in  the  definition  of  the  law  of  combining  weights 
stated  above,  for  in  that  law  no  particular  scale,  or  way  of  fixing 
values  for  the  numbers,  is  either  given  or  required.  Two  other  terms, 
namely,  equivalent  weights  (below),  and  atomic  weights  (next  sec- 
tion), however,  do  refer  to  a  fixed  scale  and  to  a  definite  numerical 
value  for  each  element. 

The  equivalent  weight  of  each  element  is  that  weight  which 
combines  with  8  parts  of  oxygen  or  1.008  parts  of  hydrogen.  The 
equivalent  weights  of  the  different  elements  (e.g.,  copper  31.8,  mer- 
cury 100.3,  chlorine  35.46,  etc.)  are  so  designated,  because  they  are 
equivalent  to  the  extent  that  they  combine  with  equal  amounts  of 
oxygen  (or  of  hydrogen).  These  numbers  are  frequently  used  in 
chemical  work. 

Atomic  Weights.  —  The  chemist  frequently  uses  the  idea  of 
equivalents  and  the  values  (p.  60)  we  have  given  them.  But  far 
more  often  he  employs  a  slightly  differing  set  of  numbers,  which  he 
calls  atomic  weights.  The  following  list  shows  the  elements  whose 
equivalents  we  have  been  discussing,  along  with  one  or  two  others, 
added  by  way  of  furnishing  a  fair  sample,  and  gives  both  sets  of 
weights  for  the  purpose  of  comparison: 


ELEMENT. 

EQUIVA- 
LENT 
WEIGHT. 

ATOMIC 
WEIGHT. 

ELEMENT. 

EQUIVA- 
LENT 
WEIGHT. 

ATOMIC 
WEIGHT. 

Oxygen     .    .    . 

8 

16 

Iron       .... 

27.92 

55.84 

Copper     .    .    . 
Sulphur    .    .    . 

31.8 
16.03 

63.6 
32.06 

Magnesium     . 
Carbon     .    .    . 

12.16 
3.00 

24.32 
12.00 

Mercury      .    . 

100.3 

200.6 

Aluminium 

9.03 

27.1 

Chlorine  .    .    . 

35.46 

35.46 

Sodium    .    .    . 

23.00 

23.00 

Hydrogen    .    . 

1.008 

1.008 

Bromine       .    . 

79.92 

79.92 

64  INORGANIC  CHEMISTRY 

It  will  be  seen  that  some  equivalents  have  been  multiplied  by  two, 
the  first  four  and  those  of  iron  and  magnesium,  for  example;  some 
have  been  multiplied  by  three,  like  that  of  aluminium;  some  by  four, 
like  that  of  carbon;  and  some  remain  unchanged,  like  those  of  chlo- 
rine, hydrogen,  sodium,  and  bromine. 

Explanation  of  the  Law  of  Combining  Weights:  Atoms 
and  Molecules.  —  To  explain  the  law  of  combining  weights  it  was 
found  necessary  to  use  the  third  kind  of  explanation  (p.  11),  namely, 
the  making  of  an  hypothesis.  The  details  of  how  two  substances 
combine  cannot  be  seen,  so  chemists  had  to  imagine  some  details 
which  would  account  for  the  possession  of  an  individual  unit  weight 
by  each  element.  If  oxygen,  for  example,  is  composed  of  minute, 
invisible  particles,  which  are  all  alike  in  weight,  and  hydrogen,  sul- 
phur, and  the  other  elements  are  of  the  same  nature,  except  that  the 
weight  of  the  particle  of  each  kind  of  element  is  different,  we  have  the 
basis  of  an  explanation.  We  have  to  suppose,  further,  that,  when 
elements  combine,  the  individual  particles  adhere  in  pairs  or  groups, 
as  wholes,  and  are  never  broken.  In  this  way  the  particle  of  each 
variety  of  elementary  matter  will  have  a  definite,  unchangeable 
weight,  which  will  be  one  of  its  fixed  properties.  If  the  relative 
weights  of  the  particles  of  oxygen,  copper,  and  hydrogen  are  in  the 
proportion  of  the  numbers  in  the  table,  namely,  16  :  63.6  :  1.008,  the 
whole  situation  becomes  clear.  Chemical  union  must  consist,  in 
detail,  in  the  union  of  the  particles  of  the  elements  to  form  the  par- 
ticles of  the  compound.  For  each  particle  of  cupric  oxide  (p.  57)  the 
proportion  31.8  :  8,  or  one  particle  of  copper  (63.6)  to  one  particle  of 
oxygen  (16),  is  required. 

For  each  particle  of  water,  where  the  proportion  of  oxygen  to 
hydrogen  is  16  :  2.016  (or  8  :  1.008),  evidently  one  particle  of  oxygen 
and  two  particles  of  hydrogen  are  necessary.  Varying,  intermediate 
proportions  are  impossible,  because  the  particles  of  the  elements  are 
permanent,  are  never  broken,  and  combine  as  wholes,  and  in  a  uni- 
form way  through  the  mass.  The  only  possible  variation  would  be 
to  take  different  relative  numbers  of  the  particles  —  for  example, 
two  of  oxygen  to  two  of  hydrogen  (2X16  :  2X1. 008).  But  this 
product  would  have  a  different  composition  from  water,  and  would  not 
be  water.  This  compound,  with  the  double  proportion  of  oxygen, 
is  indeed  known  (it  is  hydrogen  peroxide),  and  is  the  only  other 
known  compound  of  these  two  elements. 

This  theory  fully  explains  why  the  combining  proportions  of 


COMBINING   PROPORTIONS  BY  WEIGHT  65 

each  element,  in  different  compounds,  can  always  be  expressed  by  a 
fixed,  unit  number  (which  represents  the  weight  of  the  ultimate 
particle  of  that  element),  multiplied,  when  necessary,  by  a  whole 
number  (representing  the  number  of  particles  of  the  element  required 
to  form  a  particle  of  the  compound  in  question). 

This  explanation  was  first  offered  by  Dal  ton,  a  schoolmaster  of 
Manchester,  in  1802.  Borrowing  an  idea  from  the  speculations  of 
the  Greek  philosophers,  he  called  the  particles  of  elements  atoms 
(Gk.  aro/Aos,  not  cut,  or  not  divided).  The  atoms  of  any  one  element 
are  all  alike  in  weight,  as  well  as  in  other  properties,  but  the  atoms 
of  different  elements  differ  in  weight. 

The  particles  made  by  uniting  two  or  more  atoms,  as  in  forming 
a  particle  of  a  compound,  are  called  molecules  (Dim.  of  Lat.  moles, 
a  mass).  . 

The  chemical  combination  of  two  simple  substances  consists, 
then,  in  an  elaborate  re-grouping  of  the  atoms  of  both  elements  so 
that  molecules  of  the  compound  are  formed.  Definite  proportions 
by  weight  are  required,  in  order  that  the  atoms  of  each  element  may 
be  available  in  the  correct  proportion,  1  atom  :  1  atom  or  1  :  2  or 
2  :  3,  or  in  some  similar,  usually  simple  ratio. 

The  result  was  called  the  atomic  theory.  For  long  it  remained 
an  hypothesis.  Recently,  however,  we  have  obtained  independent 
proof  that  molecules  and  atoms  are  real,  for  we  can  now  count  and 
weigh  individual  molecules,  and  we  even  know  something  of  the 
inside  structure  of  atoms. 

The  fundamental  numbers,  one  for  each  element,  being  the 
relative  weights  of  the  atoms,  are  called  atomic  weights. 

Atomic  Weights  Again.  —  It  was  noted  above  (p.  63)  that  the 
atomic  weights  are  often  multiples  of  the  equivalents  by  whole 
numbers.  Thus,  the  equivalent  of  aluminium  is  9.03,  and  the  atomic 
weight  27.1;  the  equivalent  of  oxygen  8  and  the  atomic  weight  16. 
This  multiplication  was  made  to  obtain  the  true  relative  weights  of 
the  atoms.  No  facts  discussed  in  this  chapter  can  enable  us  to 
decide  upon  the  relative  weights  of  the  atoms.  In  a  later  chapter, 
however  (Chap.  XII),  additional  facts  will  be  encountered  which 
enable  us  to  reach  this  decision.  The  reader,  must,  therefore,  for 
the  present  accept  the  atomic  weights  and  use  them,  pending  the 
presentation  of  proofs  that  they  are  correct.  The  step  from  equiva- 
lents to  atomic  weights  is  taken  before  the  justification  of  it  can  be 
given,  because  otherwise  formulae  (see  next  chapter),  which  are 


66  INORGANIC  CHEMISTRY 

based  on  atomic  weights,  could  not  be  used  in  the  earlier  chapters, 
and  so  the  advantages  their  employment  offers  would  be  sacrificed. 

Attention  may  be  called  to  one  fact  which  shows  that  the  data 
presented  thus  far  do  not  permit  us  to  decide  upon  the  best  final  unit 
weights.  In  the  table  on  p.  59,  we  find  two  different  compounds  of 
oxygen  and  copper  in  which  the  proportions  are  8  :  31.8  and  8  :  63.6. 
Similarly,  8  parts  of  oxygen  combine  with  100.3  and  with  200.6  parts 
of  mercury;  and  1.008  of  hydrogen  with  8  and  with  16  parts  of  oxy-  - 
gen.  In  each  case,  the  two  compounds  are  equally  important,  and 
there  is,  therefore,  no  basis  for  deciding  whether  to  select  the  number 
31.8  or  63.6  for  the  unit  weight  of  copper,  100.3  or  200.6  for  mercury, 
8  or  16  for  oxygen.  Evidently  some  other  kind  of  information  (see 
Chap.  XII)  is  required  to  enable  us  to  make  the  decision. 

A  little  thought  will  show  that  the  atomic  weights  have  all  the 
properties  which  we  have  shown  to  belong  to  the  equivalent  weights. 
The  atomic  weight  is  the  unit  of  weight  (p.  65)  actually  used  in 
expressing  the  proportions  of  each  element  in  all  its  compounds. 
The  integral  factors  are,  of  course,  different  from  those  which  would 
be  employed  in  expressing  the  composition  of  the  same  substance  in 
terms  of  equivalents,  because  many  of  the  latter  have  been  mul- 
tiplied by  small  integers  already  in  course  of  being  made  into  atomic 
weights.  But  the  multiplication  has  in  every  case  been  by  an  in- 
teger, so  that  the  new  numbers  are  just  as  serviceable  as  are  the  old 
ones  (p,  62). 

To  the  reasons  given  above  for  the  choice  of  oxygen  as  the  funda- 
mental element,  and  the  value  8  for  its  equivalent  weight,  one  other 
may  now  be  added.  The  majority  of  the  atomic  weights,  calculated 
on  this  basis  from  the  experimental  results,  fall  so  close  to  being 
integers  that  the  nearest  round  numbers  are  exact  enough  for  ordinary- 
use.  Thus  in  the  above  list  seven  of  the  twelve  atomic  weights  are 
within  0.1  of  the  nearest  whole  number.  This  convenience  disappears 
when,  for  example,  hydrogen  with  the  value  1  (instead  of  1.008)  is 
made  the  basis. 

As  we  have  seen,  the  chemist  does  not  use  a  single  unit  of  weight 
(the  gram),  as  does  the  physicist.  He  employs  a  different  unit  of 
weight  (the  atomic  weight)  for  each  of  the  eighty  elements.  This 
does  not  represent  an  arbitrary  decision  of  the  chemist,  however.  It 
is  due  to  the  fact  that  the  atoms  of  any  one  element  have  the  same 
weight,  but  that  the  atoms  of  different  elements  have  different 
weights.  The  atom  of  uranium  is  238  times  as  heavy  as  that  of 
hydrogen,  and  its  combining  proportions,  therefore,  are  in  general 


COMBINING  PROPORTIONS  BY  WEIGHT  67 

greater  in  the  same  ratio,  while  the  atoms  of  the  other  elements  have 
weights  falling  between  these  limits. 

A  complete  list  of  atomic  weights  is  printed  on  the  inside  of  the 
cover  at  the  back  of  this  book. 

Summary.  —  In  this  chapter  we  have  encountered  three  addi- 
tional facts  which  are  characteristics  of  chemistry.  This  chapter 
adds  an  important  item  to  our  statement  of  the  scope  of  the  science 
(c/.  p.  54),  which,  therefore,  now  reads  as  follows:  Chemistry  deals 
with  the  quantitative  study  of  the  changes  in  composition  and  con- 
stitution which  substances  undergo  and  with  the  transformations  of 
energy  which  accompany  them.  To  express  the  quantitative  rela- 
tions which  are  observed,  a  different  unit  of  weight  is  employed  for 
each  element,  and  is  known  as  the  atomic  weight  of  the  element. 

There  are  other  important  characteristics  of  chemical  phenomena, 
mostly  concerned  with  the  conditions  (p.  40),  but  the  six  which  have 
been  given  are  sufficient,  for  the  present,  to  guide  us  in  the  systematic 
study  of  the  behavior  of  the  elements  and  their  chief  compounds. 

It  may  not  be  out  of  place  to  indicate  which  are  the  most  important  condi- 
tions. 

The  first  condition  whose  influence  we  are  likely  to  notice  in  chemical  work  is 
that  of  temperature.  The  accelerating  effect  of  rise  in  temperature  on  the  speed  of 
all  chemical  changes  (see  Chap.  V),  and  van't  Hoff's  law  (q.v.)  in  regard  to  the 
effect  of  temperature  on  the  direction  of  chemical  change,  describe  the  most  im- 
portant characteristics  of  this  influence. 

The  second  condition  whose  effects  we  continually  observe  is  that  of  concen- 
tration. This,  and  not  chemical  affinity,  as  many  suppose,  determines  chemical 
behavior  in  the  majority  of  familiar  actions.  It  is  described  by  the  law  of  con- 
centration (q.v.),  or  "mass  action,"  as  it  is  often  inappropriately  called  to  the  great 
detriment  of  clearness.  Brin's  method  of  obtaining  oxygen  furnishes  the  first 
conspicuous  case  of  the  influence  of  this  condition  which  we  shall  encounter.  If 
this  and  many  other  examples  are  passed  over  without  discussion,  it  is  only 
because  we  must  wait  until  much  chemical  experience  has  been  gained  before  this 
principle  can  be  understood.  Pressure  is  the  familiar  measure  of  concentration 
in  gases. 

A  third  condition  of  great  importance  in  many  —  perhaps  most  —  chemical 
actions  is  the  presence  of  a  catalytic  or  contact  agent  (q.v.). 

Exercises.  —  1.  For  the  purpose  of  recording  the  results  of 
quantitative  experiments,  why  do  we  prefer  percentages  (p.  56)  to  the 
actual  weights  themselves? 

2.   To  test,  the  correctness  of  the  statements  on  p.  61,  take  mercury 


68  INORGANIC  CHEMISTRY 

as  the  basal  element  and  250  as  its  combining  weight,  and  work  out 
from  the  data  on  pp.  59  and  60  the  corresponding  combining  weights 
of  the  other  five  elements.  Then  show  that  the  values  obtained  have 
the  same  property  as  have  the  equivalents. 

3.  Express  in  terms  of  atomic  weights,  or  their  integral  multiples, 
the   composition  of  cupric  oxide,   cupric  chloride,   sulphur  mono- 
chloride. 

4.  Show  that  doubling  the  atomic  weight  Oi  chlorine  would  give 
an  available  combining  number. 


CHAPTER  IV 

SYMBOLS,   FORMULA,   EQUATIONS,   CALCULATIONS 

A  CONSIDERATION  of  the  contents  of  the  foregoing  chapter  will  show 
that  the  complete  description  of  a  chemical  change  must  be  exceed- 
ingly involved.  In  a  moderately  complex  action,  such  as  that  of 
sodium  chloride  upon  silver  nitrate  (p.  20),  we  should  say  that  sodium 
chloride,  composed  of  one  atomic  weight  each  of  sodium  and  chlorine, 
when  brought  in  contact  with  silver  nitrate,  composed  of  one  atomic 
weight  each  of  silver  and  nitrogen  and  three  atomic  weights  of  oxygen, 
gave  silver  chloride,  composed  of  one  atomic  weight  each  of  silver  and 
chlorine,  and  sodium  nitrate,  composed  of  one  atomic  weight  each  of 
sodium  and  nitrogen  and  three  atomic  weights  of  oxygen.  Such  a 
statement,  while  it  would  give,  all  the  facts  in  the  quantitative  point 
of  view,  would  be  difficult  to  grasp  and  lacking  in  perspicuity. 

Symbols  and  Formulse.  —  In  order  to  represent  the  nature  of 
a  chemical  change  in  a  form  which  may  be  taken  in  at  a  glance,  the 
chemist  is  in  the  habit  of  using  certain  symbols,  first  introduced  by 
Berzelius.  Thus,  the  letters  Ag  represent  one  atomic  weight  (i.e., 
107.88  parts)  of  silver  (argentum),  and  0  represents  one  atomic  weight 
(i.e.,  16  parts)  of  oxygen.  In  other  words,  the  symbol  of  an  element 
means  one  chemical  unit  weight  (atomic  weight)  of  the  element. 
Since  the  names  of  many  elements  begin  with  the  same  initial,  two 
letters  have  frequently  to  be  used  to  distinguish  them.  Thus,  C 
stands  for  one  atomic  weight  (12  parts)  of  carbon,  Ca  for  one  atomic 
weight  (40.07  parts)  of  calcium,  Cl  for  35.46  parts  of  chlorine.  When 
the  names  of  the  elements  are  not  the  same  in  all  languages,  resort  is 
frequently  had  to  Latin.  Thus,  Cu  stands  for  one  weight  of  copper 
(cuprum),  Fe  is  used  for  iron  (ferrum),  Hg  for  mercury  (hydrargyrum). 
From  German  we  have  Na  for  sodium  (natrium)  and  K  for  potassium 
(kalium) . 

To  represent  a  compound,  the  symbols  of  the  elements  which  it 
contains  are  placed  side  by  side,  small  numbers  indicating  multiples 
of  the  atomic  weights  where  they  occur.  Thus,  sodium  chloride  is 
represented  by  the  symbols  NaCl  (=  23  of  sodium  to  35.46  of  chlo- 


70  INORGANIC  CHEMISTRY 

rine),  silver  nitrate  by  the  symbols  AgN03  (=  107.88  ol  silver,  14.00 
of  nitrogen,  and  3  X  16  of  oxygen).  Such  a  combination  of  symbols 
is  called  a  formula.  The  symbols  composing  a  formula,  taken  by 
themselves,  do  not  stand  for  any  definite  quantity;  each  is  one 
factor  of  a  proportion.  Ag  means  the  proportion  of  107.88  parts  of 
silver  to  the  proportions  of  the  other  elements  represented  by  the  other 
symbols  which  may  be  connected  with  it.  The  symbols  are  inter- 
national (see  list  inside  rear  cover). 

Equations.  —  It  is  now  possible  to  abbreviate  the  condensed 
statements  we  have  been  using  to  represent  the  substances  and  their 
quantities  in  chemical  reactions.  Thus,  the  statement  on  p.  13, 
when  translated  into  symbols,  is  as  follows : 

Iron  (2  X  55.84)  +  Oxygen  (3  X  16)  -»  Ferric  oxide  (159.68) 
or:      .  2Fe  +  30 -»  Fe2O3 

that  on  p.  16  is:  Fe  +  S  ->  FeS. 

The  chemical  action  above  mentioned  (p.  20)  appears  as  follows: 
NaCl  +  AgN03  -»  AgCl  +  NaNO3. 

This  expression  contains  all  that  was  conveyed  by  the  words  which 
were  written  out  in  full.  The  arrow  indicates  that  the  materials  on 
the  left-hand  side  pass,  in  the  chemical  transformation,  into  those 
on  the  right-hand  side.  Such  symbolic  expressions  are  called  equa- 
tions. 

It  will  be  observed  that,  in  the  first  equation,  we  employ  the  form 
2Fe  before  combination  and  Fe2  (in  Fe2O3)  after  it.  The  reasons  for 
this  usage  will  become  clear  as  we  proceed.  We  note  simply  that 
2Fe  means  2  separate  atomic  weights  of  iron,  as  3Fe203  would  mean 
three  separate  formula-weights  of  oxide  of  iron.  The  same  substance, 
iron,  might  appear  as  5Fe  or  8Fe  in  other  equations,  according  to  the 
proportion  needed.  But  Fe2O3  is  a  group  of  five  atomic  weights 
united  chemically.  The  substance  ferric  oxide  never  contains  either 
•more  or  less  than  two  atomic  weights  of  the  element  iron,  and  its 
formula  is  invariable.  Thus  the  regular  integers  multiplying  the 
atomic  weights  in  the  composition  of  a  particular  compound  are 
written  after  the  symbols  of  the  elements,  while  more  arbitrary 
coefficients  which  change  from  one  use  of  the  substance  to  another  are 
written  in  front.  When  no  coefficient  appears  in  front  of  a  symbol  or 
formula,  1  is  to  be  understood. 


SYMBOLS,  FORMULA,  EQUATIONS,  CALCULATIONS          71 

Much  practice  is  required  to  enable  one  to  make  and  understand 
equations.  The  reader  should,  therefore,  at  once  turn  back  to  the 
statements  on  pp.  13,  16,  17,  18,  obtain  the  necessary  atomic 
weights  and  symbols  from  the  table  at  the  end  of  the  book,  and  con- 
struct the  equation  in  each  case. 

We  shall  find  later  that  there  are  two  substances  containing  nothing  but 
oxygen,  and  that  each  is  a  compound  of  the  element  with  itself.  The  molecular 
formulae  of  these  two  are  O2  (oxygen)  and  O3  (ozone).  Thus  O  or  2O  or  3O 
would  all  be  used  for  different  proportions  of  the  substance  O,  if  such  a  substance 
were  known,  and  O  would  be  used  for  a  substance  made  of  oxygen,  but  different 
from  oxygen  or  ozone.  Molecular  formulae  will  not  be  employed  here  until  after 
Avogadro's  law  has  been  discussed.  They  will  then  be  used  exclusively. 

The  object  in  writing  a  series  of  formulae  in  the  above  manner  is  to  show 
that  the  system  upon  the  left-hand  side,  consisting  of  certain  substances  whose 
composition  and  properties  we  know,  is,  under  the  conditions  of  the  experiment, 
unstable,  and  changes  into  the  system  upon  the  right-hand  side,  whose  nature 
we  also  know.  The  materials  on  the  two  sides  are  essentially  different,  for  the 
transformation  represented  is  a  chemical  change.  It  is  somewhat  anomalous, 
therefore,  that,  to  connect  two  sets  of  things  which  are  essentially  different,  the 
sign  =  is  usually  employed.  To  call  this  a  chemical  equation  is  still  more  anoma- 
lous, since  it  is  precisely  in  the  chemical  point  of  view  that  the  difference  between 
the  two  sides  is  most  strongly  to  be  emphasized.  It  represents  two  sets  of  things 
which  are  different,  and  not  alike  chemically.  The  physical  properties  of  the  two 
sets  of  substances  are  likewise  totally  unlike.  There  is  only  one  respect  in  which 
the  materials  on  the  two  sides  agree,  and  that  is  that  their  mass  is  not  different. 
This  is,  however,  merely  an  example  of  the  law  of  conservation  of  matter,  and 
need  not,  therefore,  be  specially  commemorated  in  the  form  in  which  we  write 
every  equation.  It  may  be  assumed  that  the  equality  in  mass  holds  for  all  chemi- 
cal changes  until  some  case  where  it  does  not  hold  shall  have  been  discovered. 
Above  all  it  must  be  remembered  that  the  chemical  equation  is  not  an  algebraic 
expression;  it  is  not  subject  to  the  rules  of  algebra.  It  is  a  brief  expression,  in 
terms  of  the  atomic  weights,  of  the  distribution  in  kind  and  quantity  of  the  con- 
stituents of  a  system  before  and  after  chemical  change. 

Making  Formulae.  —  To  make  the  formula  of  a  compound  sub- 
stance, assuming  the  formula  to  be  unknown,  two  kinds  of  informa- 
tion are  required.  We  ascertain  (1)  by  measurement  the  proportion 
by  weight  of  the  constituents  in  the  compound.  We  require  also 
(2)  to  know  the  chemical  unit  weights  —  the  atomic  weights  —  which 
have  been  accepted  by  chemists  for  each  constituent  element.  By 
factoring  the  terms  of  the  first  proportion  so  that  one  factor  in  each 
case  is  the  atomic  weight,  we  discover  whether  multiples  of  the  atomic 
weights  will  be  required  to  represent  the  composition  of  the  sub- 


72  INORGANIC  CHEMISTRY 

stance,  and  if  so  what  these  must  be.  An  illustration  will  make  the 
process  clear. 

Suppose  the  problem  is  to  make  the  formula  of  dried  rust.  By 
weighing  before  and  after  the  change,  we  get  the  weight  of  the  iron 
and  of  the  corresponding  amount  of  oxygen  in  the  rust  it  produces. 

If  we  took  2  g.  of  iron  we  should  get  about  2.86  g.  of  rust.     So  that 

2 

the  proportion  of  iron  to  oxygen  is  — — .     Now,  in  the  formula,  the 

u.ob 

same  ratio  must  be  represented  by  means  of  multiples  of  the  atomic 
weights  (p.  69).  We  therefore  divide  the  quantity  of  each  element 
by  the  corresponding  atomic  weight.  This  gives  us  the  factors  by 
which  the  atomic  weights  are  to  be  multiplied.  The  atomic  weights 
are  55.8  and  16  respectively:  2  -^  55.8  =  0.0358,  and  0.86  -f-  16  = 

2      V,  65.8  X  0.0358      -T 

0.0537.     The  proportion  —  then  becomes    16  Q  x  Q  Q537  •     Now 

this  proportion  must  be  capable  of  expression  in  terms  of  integral 
multiples  of  the  atomic  weights.  We  find  that  the  greatest  common 
measure  of  the  two  factors  is  0.0179.  Dividing  above  and  below  by 

this,  we  obtain  the  ratio  T-^- .     Substituting  the  symbols  for  the 

lo.O  X  o 

FP  X  2 

atomic  weights,  the  proportion  appears  as  — ,  and  the  formula 

O  X  o 

is  therefore  Fe2O3. 

Applying  the  same  process  to  cupric  oxide  (p.  56),  we  start  with 
the  result  of  the  measurement:  copper  :  oxygen  : :  85  :  21. 

Copper  =  85  =  63.57  X  1.3  _  Cu  X  1.3  _  Cu 
Oxygen      21         16X1.3      =  O  X  1.3  =  =  O   ( 

If  the  composition  of  the  substance  has  been  stated  in  percentages, 
the  same  device  is  used.  Thus,  the  case  of  sodium  sulphate  works 
out  as  follows : 


Element. 

Percentage. 

At.  Wt.     Quotient 

-7- 

Formula 

Sodium     

32  43 

23     X  1  41 

0  705 

NaX2 

Sulphur    

22  55 

32     X  0  705 

0.705 

S 

Oxygen 

45  02 

16     X2.814 

0.705 

OX4 

The  formula  is,  therefore,  Na2S04. 

It  is  obvious  that,  after  we  have  found  out  what  elements  com- 
pose a  given  compound,  we  are  still  unable  to  write  its  formula.     We 


SYMBOLS,  FORMULA,  EQUATIONS,  CALCULATIONS          73 

may  not  simply  set  the  symbols  down,  side  by  side.  A  measurement 
must  be  made,  in  order  that  we  may  find  out  the  factors  by  which  the 
atomic  weights  are  to  be  multiplied. 

Writing  Equations.  —  To  make  the  equation  representing  a 
chemical  change  we  take  the  following  steps: 

1.  We  note  the  formulas  of  the  substances  used. 

2.  We  recognize  by  their  properties  the  substances  produced,  and 
learn  their  formulae.     This  we  do,  either  by  measurement  and  calcu- 
lation, as  shown  above,  or  we  find  in  a  book  the  formulae  as  they 
have  been  determined  by  the  experimental  work  of  chemists. 

3.  We  then  write  a  skeleton  equation.     In  the  first  example  dis- 
cussed above  (p.  70),  this  is: 

Skeleton  equation:  Fe  +  0 


We  are  careful  to  place  the  substances  used  on  the  left,  and  to  point 
the  arrow  towards  those  which  are  produced. 

4.  Finally,  we  balance  the  equation,  by  placing  the  proper 
coefficients  before  the  formulae.  This  last  operation  requires  experi- 
ence for  its  rapid  performance.  A  good  rule  is  to  begin  by  picking 
out  that  one  of  the  formulae  which  contains  the  largest  number  of 
atomic  weights,  no  matter  upon  which  side  it  appears.  Here,  this 
formula  is  Fe2O3.  We  then  reason  that,  to  obtain  Fe2  we  require 
2Fe,  and  to  obtain  O3  we  require  3O,  and  accordingly  we  place  these 
coefficients  before  the  appropriate  symbols,  thus: 

Balanced:  2Fe  +  30  -»  Fe203. 

It  is  hardly  necessary  to  add  that  a  chemical  equation  gives  the 
proportions  of  the  materials  and  nothing  more.  The  physical  condi- 
tions, for  example,  whether  the  substances  are  dissolved  in  a  liquid,  or 
are  in  the  state  of  gas,  or  are  at  a  high  temperature,  have  no  place  in 
it.  The  physical  properties  of  the  substances  concerned,  and  also 
the  energy  in  the  form  of  heat  or  electricity  which  may  appear  or  dis- 
appear in  the  process,  are  likewise  left  entirely  out.  A  question  in 
regard  to  the  nature  of  a  particular  chemical  change  demands  in 
answer  a  full  statement  of  all  these  things.  The  equation  is  therefore 
an  essential  part,  but  only  a  part,  of  such  a  statement. 

That  the  formulae  and  equations  can  deal  only  with  the  material  part  of  the 
substances  undergoing  change,  and  not  with  their  energy  (p.  35),  is  shown  by  a 
moment's  consideration.  Consistency  is  to  be  secured  only  by  holding  tnat  the 


74  INORGANIC  CHEMISTRY 

symbol  S,  whether  alone  or  combined  with  others,  stands  for  the  matter-part  of 
the  sulphur  (for  the  element,  in  fact,  see  p.  23).  It  is  32  parts  by  weight  of  sul- 
phur-matter. Only  in  this  way  does  it  preserve  the  same  significance  on  both  sides 
of  the  equation  S  +  O2  — *•  SO2.  If  S  on  the  left  side  stood  for  the  free  substance 
sulphur,  then  it  would  stand  for  32  parts  of  sulphur-matter  plus  the  appropriate 
amount  of  energy.  In  this  case  the  S  on  the  right  side  would  have  a  different  sig- 
nification, and  represent  a  less  amount  of  energy.  This  is  on]y  the  beginning  of 
the  difficulty,  for  we  then  find  that  S  in  H2SO4  represents  the  same  weight  of 
sulphur-matter  with  still  another  proportion  of  energy,  and  S  has  as  many  inter- 
pretations as  there  are  formulae  in  which  it  occurs.  Clearly,  all  references  to 
energy  should  be  rigidly  excluded  from  equations,  and  thermochemical  data  can 
never  be  given  in  connection  with  them  without  complete  sacrifice  of  consistency. 
In  this  book,  however,  the  habit  of  writing  thermochemical  equations,  being 
universal,  is  frequently  followed  when  thermochemical  data  are  given. 

Units  of  Measurement  in  Chemical  Work.  —  In  chemical 
work  temperatures  are  invariably  measured  on  the  Centigrade  scale. 
The  temperature  of  a  mixture  of  ice  and  water  is  the  zero  point.  The 
temperature  of  the  steam  which  rises  from  water  boiling  under  a 
pressure  of  one  atmosphere  is  represented  by  100°.  The  interval 
between  those  two  points  is  divided  into  one  hundred  equal  parts. 

For  the  expression  of  length,  weight,  and  volume,  the  metric 
system  is  employed.  The  unit  of  this  system  is  the  meter,  which  is 
subdivided  into  decimeters,  centimeters  (cm.),  and  millimeters  (mm.). 
For  small  measurements  the  last  subdivision  is  taken  as  the  unit.  A 
cubic  centimeter  (c.c.)  is  the  unit  of  volume  for  small  measurements. 
For  larger  ones  the  liter,  which  contains  1000  cubic  centimeters,  is 
used.  The  unit  of  weight  is  that  of  one  cubic  centimeter  of  water  at 
4°,  the  temperature  of  maximum  density.  This  is  called  the  gram.* 
For  larger  amounts  of  material  the  kilogram,  which  contains  1000 
grams  (1000  g.),  is  frequently  employed.  The  meter  is  equal  to  about 
39  J  inches  in  ordinary  measures,  and  the  centimeter  is  very  nearly  f  of 
an  inch.  One  liter  is  about  ?V  of  a  cubic  foot  and  contains  61  cubic 
inches.  One  hundred  grams  is  about  3J  ounces  avoirdupois,  arid  one 
ounce  equals  28.35  grams  (see  Appendix  I). 

Calculations.  —  As  we  have  seen  (p.  69),  the  formula  repre- 
sents the  composition  of  a  substance,  using  the  atomic  weights  as  the 
units.  We  have  learned  how  the  formula  is  calculated  from  measure- 

*  In  point  of  fact,  the  gram  is  the  one-thousandth  part  of  the  weight  of  the 
standard  kilogram  kept  in  Paris.  This  differs  from  the  weight  of  1  c.c.  of  water 
at  4°  by  less  than  0.01  per  cent. 


SYMBOLS,  FORMULA,  EQUATIONS,  CALCULATIONS          75 

ments  made  in  an  experiment  (p.  72).  We  may  now  take  up  some 
of  the  ways  of  using  the  information  contained  in  a  formula. 

Composition  from  the  Formula.     Formula-Weight.  —  To 

learn  the  composition  of  a.  substance,  such  as  potassium  chlorate 
KC103,  from  its  formula,  we  look  up  the  values  of  the  atomic  weights 
(inside  rear  cover).  We  find  K  =  39.1  parts  of  potassium,  Cl  = 
35.46  parts  of  chlorine,  and  O3  =  3  X  16  or  48  parts  of  oxygen.  The 
proportions,  in  order,  are  therefore:  39.1  :  35.46  :  48. 

What  is  the  proportion  of  oxygen  to  potassium  and  chlorine, 
together?  It  is  48  :  39.1  +  35.46,  or  48  :  74.56,  or  1  :  1.55. 

We  require  a  name  for  the  sum  of  the  weights  of  the  constituents 
indicated  in  the  formula.  This  is  called  the  formula-weight.  Thus, 
for  potassium  chlorate,  it  is  39.1  -f  35.46  -f  48,  or  122.56. 

To  Find  the  Percentage  Composition.  —  In  potassium  chlo- 
rate the  proportions  are  39.1  of  potassium,  35.46  of  chlorine,  and  48 
of  oxygen  or  a  total  of  122.56.  In  one  hundred  parts,  the  potassium 

is  -^-  X  100,  or  31.9;   the  chlorine  ^jj-  X  100,  or  28.9;   and 

the  oxygen  -^-  X  100,  or  39.1. 
IAA.OO 

Stated  in  terms  of  the  rule  of  proportion,  we  have,  for  the  potas- 
sium, 122.56  :  39.1  ::  100  :  x,  where  x  is  the  percentage  of  potassium. 

Calculations  by  Use  of  Equations.  —  We  frequently  wish  to 
know  what  weight  of  a  product  can  be  obtained  from  a  given  weight 
of  the  necessary  materials,  or  how  much  material  is  required  to 
furnish  the  desired  weight  of  a  product.  For  example,  what  weight 
of  ferrous  sulphide  can  be  made  with  100  g.  of  iron?  It  is  understood 
that  the  necessary  sulphur  is  available. 

To  avoid  the  blunders  which  are  easily  made,  observe  strictly  the 
following  rules: 

1.  Write  down  the  equation: 

Fe  +  S->FeS. 

2.  Place  under  each  formula  the  weight  it  represents: 

Fe     +     S      ->    FeS. 
55.84        32.06          87.9 


76  INORGANIC  CHEMISTRY 

3.  Read  this  expanded  equation:     In  this  case  it  reads:    55.84 
parts  of  iron  combine  with  32.06  parts  of  sulphur  to  give  87.9  parts 
of  ferrous  sulphide. 

4.  Re-read  the  original  problem:     "What  weight  of  ferrous  sul- 
phide can  be  made  with  100  g.  of  iron?"     Having  done  this,  place 
the  amount  given  in  the  problem  (100  g.  of  iron)  under  the  formula 
of  the  substance  in  question.     Then  notice  what  the  problem  asks 
("  what  weight  of  ferrous  sulphide")  and  place  an  x  under  the  formula 
of  that  substance: 

Fe  +  S  ->  FeS 
55.84  32.06  87.9 
100  g.  x 

5.  Read  the  problem  as  now  tabulated:     55.84  g.  of  iron  give 
87.9  g.  of  ferrous  sulphide,  therefore  100  g.  of  iron  will  give  x  g.  of 
ferrous  sulphide. 

6.  State  the  proportion  in  this  order  (or,  see  below) : 

55.84  :  87.9  ::  100  :  x  (=  157.5  g.). 

If  the  tabulation  in  rule  4  has  been  prepared  correctly,  this  final 
statement  as  a  proportion  is  purely  mechanical.  It  will  be  noted 
that  only  two  of  the  three  quantities  given  in  the  expanded  equation 
were  actually  used. 

6a.   Alternative  method:      At  the  sixth  step,  we  may  also  say: 
If  55.84  g.  of  iron  give  87.9.g.  of  ferrous  sulphide,  1  g.  of  iron  will  give 

87.9 

'    g.  (=  1.575  g.)  of  ferrous  sulphide.     Then,  if  1  g.  of  iron  gives 

OO.O4 

1.575  g.  of  ferrous  sulphide,  the  100  g.  of  iron  will  give  100  X  1.575  g. 
(=  157.5  g.)  of  ferrous  sulphide. 

Warnings.  —  In  solving  the  exercises  at  the  end  of  the  chapter, 
beware  of  three  kinds  of  mistakes,  which  are  commonly  made. 

1.  Do  not  read  the  problem  carelessly  and  make  the  equation 
backwards,  that  is,  with  the  sides  reversed.     Focus  attention  first 
on  the  exact  chemical  change  involved. 

2.  Do  not  speak,  or  think  of  the  symbols  Fe  and  S  as  standing  for 
"1  part"  of  iron  or  sulphur.     They  stand  for  1  chemical  unit,  or 
atomic  weight,  or  atom,  in  each  case,  that  is,  for  "55.84  parts"  and 
"32.06  parts,"  respectively. 

3.  Follow  the  rules  laid  down  above.     The  chemist  follows  these 
rules.     The  beginner  always  thinks  he  can  do  without  them,  and  he 
fails  in  consequence.     Writing  the  equation  in  expanded  form  (rule 


SYMBOLS,  FORMULAE,  EQUATIONS,  CALCULATIONS  77 

2)  and  reading  the  problem  into  it  (rules  4  and  5)  are  absolutely 
essential  steps. 

Another  .  Example.  —  What  weight  of  hydrogen  is  required  to 
reduce  45  g.  of  magnetic  oxide  of  iron  to  metallic  iron? 

Following  the  rules,  as  before,  we  reach  the  expanded  equation: 

Fe3O4  +        4H2     ->       3Fe      +          4H2O. 

3  X  55.84  +  4  X  16       8  X  1.008      3  X  55.84      4(2  X  1.008  +  16) 

167.52  +  64  8.064  167.52  4  X  18.016 

231.52  8.064  167.52  72.064 

45  g.  x 

Observe  that  the  atomic  weights  are  multiplied  by  the  sub-numbers, 
so  that,  for  example,  Fe3  =  3  X  55.84.  Observe  also  that  the 
formula-weights  are  multiplied  by  the  coefficients,  when  such  occur, 
in  front  of  the  formulae,  so  that,  for  example,  4H20  =  4  X  18.016. 

The  proportion  231.52  :  8.064  ::  45  :  x  (=  1.57)  supplies  the 
answer,  1.57  grams  of  hydrogen. 

Using  the  alternative  plan:    If  231.52  g.  of  magnetic  oxide  are 

o   f\f\A 

reduced  by  8.064  g.  of  hydrogen,  1  g.  will  be  reduced  by  -     —  g. 


(  =  0.035  g.)  of  hydrogen.  Hence,  if  1  g.  of  magnetic  oxide  is  reduced 
by  0.035  g.  of  hydrogen,  45  g.  will  be  reduced  by  45  X  0.035  g. 
(=  1.57  g.)  of  hydrogen. 

Exercises.  —  L  1.   What  weight  of  mercury  is  obtained  from  120  g. 
of  mercuric  oxide? 

2.   What  weight  of  mercuric  oxide  will  furnish  20  g.  of  oxygen? 
,3.   What  weight  of  silver  chloride  is  obtained  from  50  g.  of  silver 
nitrate  (p.  70)? 
^.   What  weight  of  rust  may  be  obtained  from  10  g.  of  oxygen? 

5!    How  much  silver  is  contained  in  100  g.  of  an  impure  specimen 
of  silver  chloride  which  is  33  per  cent  sand? 

6.    If  26  g.  of  mercurous  oxide  are  required  to  give,  by  heating, 
1  g.  of  oxygen,  what  is  the  formula  of  the  substance? 
^   7.   What  are  the  formulae  of  the  substances  possessing  the  follow- 
ing percentage  compositions  (p.  72). 

I  II  III 

Magnesium,     25.57  Sodium,     32.43  Potassium,  26.585 

Chlorine,         74.43  Sulphur,     22.55  Chromium,  35.390 

Oxygen,     45.02  Oxygen,  38.025 


78  INORGANIC  CHEMISTRY 

8.  What    are    the    percentage    compositions    of    hausmannite 
Mn304,  potassium  bromide  KBr,  ferrous  sulphate  FeS04? 

9.  What  weight  of  hydrogen  is  required  to  reduce  100  g.  of  ferric 
chloride  FeCl3  to  ferrous  chloride  FeCl2  (hydrogen  chloride  HC1  is 
formed)? 

10.  Which  law  of  chemistry  permits  us  to  use  symbols  in  express- 
ing the  compositions  of  substances? 

11.  Calculate  the  formula  of  the  oxide  of  tin  formed  when  2  g. 
of  tin  give  2.54  g.  of  the  oxide. 


CHAPTER  V 
OXYGEN 

WE  begin  the  more  systematic  study  of  chemistry  with  oxygen, 
for  it  is  a  most  interesting  as  well  as  useful  substance.  It  is  the 
active  substance  in  the  air.  We  depend  upon  it  for  life,  since  in  its 
absence  we  suffocate,  for  heat,  since  wood,  coal,  and  gas  will  not  burn 
without  it,  and  even  for  light  where  oil,  gas,  or  a  candle  is  employed. 

We  wish  to  know  with  which  substances  we  use  in  the  laboratory 
it  can  combine,  as  well  as  the  substances  on  which  it  has  no  action. 
This  information  will  show  us  how  to  work,  in  future,  without  inter- 
ference from  the  oxygen  in  the  air  and  whether  oxygen  has  probably 
played  a  part  in  some  experiment  or  not. 

We  take  up,  then,  (1)  the  history  of  the  element,  (2)  what  mate- 
rials contain  oxygen  (occurrence),  (3)  how  we  can  obtain  it  in  a 
pure  state  (preparation),  (4)  what  its  specific  physical  properties  as  a 
substance  are,  and  (5)  what  it  does,  and  what  it  can  not  do  in  nature 
and  in  the  laboratory  (chemical  properties).  The  classification  of 
the  facts  about  this,  and  other  substances,  under  five  heads  is  some- 
what mechanical,  but  has  the  advantage  of  enabling  the  reader 
quickly  to  find  any  required  information. 

History  of  Oxygen.  —  While  many  elements  which  are  less  easily 
obtainable  than  oxygen  have  been  recognized  as  distinct  substances 
for  many  centuries,  oxygen  did  not  attain  this  position  until  the  end 
of  the  eighteenth  century.  The  reason  of  this  was  that  gases  are 
not  so  easy  to  handle  and  distinguish  as  are  solids  or  liquids,  and 
consequently  very  slow  progress  was  made  in  the  study  of  them. 

The  Chinese,  in  or  before  the  eighth  century,  knew  that  there 
were  two  components  in  the  air,  and  that  the  active  one,  yin,  com- 
bined with  some  metals,  and  with  burning  sulphur,  and  charcoal. 
They  even  knew  that  it  could  be  obtained  in  pure  form  by  heating 
certain  minerals,  of  which  one  was  saltpeter.  Leonardo  da  Vinci 
(1451-1519)  was  the  first  European  to  state  that  the  air  contained 
two  gases.  Mayow  (1669)  measured  the  proportion  of  oxygen  in 
the  air  and  discussed  fully  its  uses  in  combustion,  rusting,  vinegar- 

79 


80  INORGANIC  CHEMISTRY 

making,  and  respiration,  but  did  not  make  a  pure  sample.  Hales 
(1731)  made  it  by  heating  saltpeter,  and  measured  the  amount 
obtainable,  but  did  not  see  any  connection  between  it  and  the  air! 
Bayen  (April,  1774)  was  the  first  to  make  it  by  heating  mercuric 
oxide.  Priestley  *  was  particularly  interested  in  examining  the  nature 

^  of  the  gases  which  were  evolved  by  some  materials 

when  heated.  His  plan  was  to  fill  an  elongated  glass 
vessel  with  mercury  (Fig.  25);  to  invert  this  in  a 
trough  filled  with  the  same  metal,  and,  after  allow- 
ing the  substance  under  examination  to  float  up  into 

^ v       j the  top  of  the  tube  above  the  mercury,  to  expose  it 

[^  "^  to  the  rays  of  the  sun  concentrated  by  a  large  burn- 

ing lens.  Priestley  found  (Aug.  1,  1774)  that  one 
material,  then  known  as  "mercurius  caldnatus  per  se" 
(mercuric  oxide),  gave  off  an  unusual  amount  of  a  gas,  or  "air"  as 
he  called  it.  Quite  purposelessly,  as  he  admits,  he  thrust  a  lighted 
candle  into  it  and  was  delighted  with  the  extreme  brilliance  of  the 
flame.  He  had,  however,  no  notion  until  a  year  later  that  it  was  a 
component  of  the  air.  Even  then,  he  thought  it  was  a  compound  of 
nitric  acid,  earth,  and  phlogiston!  Scheele,  a  Swedish  apothecary, 
had  made  it  in  1771-2  from  no  less  than  seven  different  substances 
and  understood  clearly  that  atmospheric  oxygen  combined  with 
metals,  phosphorus,  hydrogen,  linseed  oil,  and  many  other  substances. 
But  his  book  was  not  published  until  1777, 
and  Priestley  is  usually  credited  with  the 
"discovery"  of  the  element! 

Lavoisier  at  first  (1773-5)  thought  that 
air  was  composed  of  nitrogen  and  "fixed  air" 
(carbon  dioxide).  Although  Priestley  had 
dined  with  him  (Oct.,  1774),  and  commu- 
nicated his  discovery,  so  far  as  he  understood  FlG  26. 
it  himself,  yet  when  Lavoisier  heated  mercuric 

oxide  in  March,  1775,  he  expected  to  obtain  "fixed  air."  Later  in 
the  same  year,  Lavoisier  held  that  air  contained  no  "fixed  air,"  and 
was  made  up  of  a  single  gas.  It  was  not  until  1777  that  he  heated 
the  metal  mercury  in  a  retort  (Fig.  26),  the  neck  of  which  projected 
into  a  bell-jar  standing  in  a  larger  dish  of  mercury.  The  air,  thus 
enclosed  within  the  jar  and  the  retort,  during  twelve  days  lost  one- 

*  An  English  nonconformist  minister  who  occupied  his  leisure  time  with 
experiments  in  chemistry,  He  afterwards  moved  to  the  United  States,  and  died 
in  Northumberland,  Pa, 


OXYGEN  81 

fifth  of  its  volume.  Simultaneously,  red  particles  of  mercuric  oxide 
accumulated  on  the  surface  of  the  mercury  in  the  retort.  The 
residual  gas  no  longer  supported  life  or  combustion,  and  hence  was 
named  by  Lavoisier  " azote"  (Gk.  d,  priv.  and  £a»},  life).  In  English 
it  is  called  nitrogen.  The  oxide,  on  being  heated  more  strongly 
by  itself,  gave  off  a  gas,  the  volume  of  which  exactly  corresponded 
with  the  shrinkage  undergone  by  the  enclosed  air,  and  this  gas  pos- 
sessed in  an  exaggerated  degree  the  properties  which  the  air  had  lost. 
The  proof  that  oxygen  was  a  component  of  the  atmosphere  was 
therefore  complete.  Lavoisier,  in  the  mistaken  belief  that  the  new 
element  was  an  essential  constituent  of  all  sour  substances,  named 
it  oxygen,  or  acid-producer  (Gk.  6£vs,  an  acid,  yewav,  to  produce). 
Cavendish  pointed  out  almost  immediately  that  there  were  sour- 
tasting  substances  which  contained  no  oxygen,  so  that  the  name 
has  no  longer  any  significance.  It  may  here  be  noted  that  hydrogen 
is  the  only  element  which  is  found  in  all  acids. 

Occurrence.  —  As  we  have  seen  (p.  24),  nearly  50  per  cent  of 
terrestrial  matter  is  oxygen.  Water  contains  about  89  per  cent,  the 
human  body  over  60  per  cent,  and  common  materials  like  sandstone, 
limestone,  brick,  and  mortar  more  than  50  per  cent  of  this  element 
in  combination.  One-fifth  by  volume  (nearly  one-fourth  by  weight) 
of  the  air  is  free  oxygen. 

Preparation  of  Oxygen.  —  1.  The  oxygen  of  commerce  is  now 
made  chiefly  from  liquefied  air  (q.v.).  The  liquid  oxygen 
boils  at  —182.4°,  but  the  nitrogen  boils  at  an  even 
lower  temperature  (  —  194°).  As  the  liquid  air  has  a 
temperature  of  about  —190°,  somewhat  above  that  of 
boiling  nitrogen,  the  latter  evaporates  much  more  freely 
than  the  oxygen.  After  a  time,  when  the  remaining 
liquid  is  almost  pure  oxygen  (96%),  the  gas  coming  off 
is  compressed  by  pumps  (100-150  atmos.)  into  the  steel 
cylinders  (Fig.  27)  in  which  it  is  sold.  In  medicine, 
patients  suffering  from  pneumonia  or  suffocation  obtain 
some  relief  by  inhaling  it  in  this  form.  It  is  also  used 
in  feeding  flames,  instead  of  air,  when  intense  heat  is 
required  (see  acetylene  torch  and  calcium  light). 

2.  Unfortunately,  it  is  difficult  to  liberate  oxygen 
from  natural  substances.  Saltpeter  (potassium  nitrate 
KNOs),  for  example,  which  is  found  in  many  soils,  and  can  be 


82  INORGANIC  CHEMISTRY 

dissolved  out  with  water,  gives  off  oxygen  (p.  79)  only  when  raised 
to  a  bright  red  heat  by  the  Bunsen  flame  or  blast  lamp.  But,  even 
at  this  temperature,  it  gives  up  only  one-third  of  the  oxygen  it  con- 
tains, leaving  potassium  nitrite  KNO2: 

KNO3  -*  KN02  +  0. 

The  mineral  pyrolusite  (manganese  dioxide  MnO2),  employed 
by  Scheele,  requires  a  higher  temperature.  It  usually  contains 
the  elements  of  water,  also,  and  gives  off  water  vapor  at  the  same 
time.  An  oxide  of  the  composition  MriaC^  remains: 


Skeleton:  Mn02  ->  Mn304  +  0. 

Balanced:  3Mn02  -*  Mn304  +  2O. 

3.  In  practice,  we  are  compelled  to  use  manufactured  sub- 
stances. Amongst  these  are  mercuric  oxide,  expensive,  but  his- 
torically interesting  (p.  80);  barium  peroxide,  formerly  used  in 
manufacturing  oxygen  on  a  large  scale  (Brin's  process);  potassium 
chlorate  (see  below),  the  most  convenient  for  laboratory  use;  and 
sodium  peroxide.  Many  other  substances  of  this  class  will  be 
encountered  in  the  sequel. 

Brin's  Oxygen  Process.  —  This  starts  from  barium  oxide  (q.v.). 
Barium  oxide  BaO  closely  resembles  quicklime  CaO,  but  differs  from 
this  substance  in  the  fact  that,  when  heated  in  air  to  about  500°,  it 
rapidly  acquires  additional  oxygen  and  gives  barium  peroxide  BaO2. 
When  barium  peroxide  is  raised  to  1000°,  this  extra  oxygen  is  given 
up  again.  Barium  oxide  contains  one  chemical  unit  weight  each  of 
the  two  constituents  and  takes  up  another  unit  of  oxygen,  so  that 
the  equation  for  the  primary  action  is: 

BaO  +  O  -»  Ba02. 

The  subsequent  decomposition  of  the  peroxide,  during  the  stage  in 
which  the  oxygen  is  made,  is  the  exact  opposite:  BaO2  —  >  BaO  +  O.* 
The  commercial  advantage  of  the  method  lies  in  the  fact  that  the 
barium  oxide,  remaining  after  the  second  stage,  can  be  used  over  and 
over  again.  This,  as  will  be  seen,  is  in  reality  a  chemical  method  of 
obtaining  oxygen  from  the  air. 

*  In  cases  where  an  action  is  reversible,  and  the  direction  depends  on  condi- 
tions which  may  be  altered,  we  write  both  equations  in  one: 

BaO  +  O  «=±  BaO2. 


OXYGEN  83 

In  practice  an  improvement  on  the  above  principle  makes  working 
more  economical.  It  is  found  that  if  the  barium  oxide  is  maintained 
at  a  temperature  of  700°,  intermediate  between  the  two  just  men- 
tioned, oxygen  is  absorbed  when  air  is  forced  under  pressure  into  the 
tubes  containing  the  oxide.  A  valve  at  the  extremity  of  the  tubes 
permits  the  escape  of  the  nitrogen.  When  the  combination  with 
oxygen  is  completed,  the  pumping  apparatus  is  reversed  and,  a  partial 
vacuum  being  created,  the  oxygen  in  combination  is  given  off  without 
any  alteration  in  temperature  being  necessary.  Thus  a  great  waste 
of  fuel  is  avoided,  and  the  process  is  rendered  more  nearly  continuous. 

Oxygen  from  Potassium  Chlorate.  —  This  is  a  white  crystal- 
line substance  used  in  large  quantities  in  the  manufacture  of  matches 
and  fireworks.  When  heated  in  a  tube  similar  to  that  in  Fig.  9,  it 
first  melts  (351°)  and  then,  on  being  more  strongly  heated,  it  effer- 
vesces and  gives  off  a  very  large  volume  of  oxygen.  Examination 
shows  that  the  whole  of  the  oxygen  it  contains  can  be  driven  out. 
The  white  material  which  remains  after  the  heating  is  potassium 
chloride  and,  when  decomposed,  it  yields  potassium  and  chlorine  in 
the  exact  ratio  of  their  atomic  weights.  Its  formula  is  thus  KC1. 
We  may  infer,  therefore,  that  the  composition  of  the  original  sub- 
stance will  be  representable  by  the  formula  KC10X,  where  x  is  the 
number  of  atomic  weights  of  oxygen.  Measurement  and  calculation 
show  x  =  3.  The  formula  is  therefore  KC103,  and  the  equation 
for  the  decomposition  (see,  however,  under  Perchlorates) : 

KC1O3  -+  KC1  +  3O. 

To  learn  the  value  of  x,  we  ascertain  the  loss  in  weight  ( =  oxygen)  which  a 
known  quantity  of  potassium  chlorate  sustains  when  heated  in  a  hard  glass  tube 
closed  at  one  end.  By  subtraction  we  get  the  weight  of  potassium  chloride  for- 
merly combined  with  the  oxygen.  In  an  actual  experiment,  2.998  g.  of  potassium 
chlorate  gave  1.169  g.  of  oxygen  and  left  1.829  g.  of  the  chloride.  The  atomic 
weights  of  potassium  and  chlorine  are  39.1  and  35.46  respectively,  and  the  for- 
mula-weight of  the  chloride  is  therefore  74.56.  Dividing  the  measured  weights 
of  oxygen  and  potassium  chloride  by  the  corresponding  atomic  and  formula- 
weights  (c/.  p.  72),  respectively:  1.169  -r-  16  =  0.07306  and  1.829  -i-  74.56  = 
0.02452.  We  observe  that  the  ratio  of  the  quotients  is  2.98  :  1,  or  almost  exactly 
3  :  1.  The  formula  is  therefore  O  X  3,  (KC1)  X  1,  or  KC1O3. 

A  peculiarity  of  this  action  is  that  admixture  of  manganese  dioxide 
increases  very  markedly  the  speed  with  which  the  decomposition  of 
the  potassium  chlorate  takes  place.  Hence,  in  its  presence,  and  it  is 


84 


INORGANIC  CHEMISTRY 


generally  mixed  with  the  chlorate  in  laboratory  experiments  (Fig.  2S). 
a  sufficient  stream  of  the  gas  is  obtained  at  a  relatively  low  tempera- 
ture (below  200°,  see  p.  97).  Hales  (p.  80)  was  the  first  to  collect  a 
gas  over  water,  in  order  that  it  might  be  kept  unmixed  with  air. 


FIG.  28. 

Familiarity  with  Physics  Required  in  the  Study  of  Chem- 
istry. —  In  mentioning  chemical  phenomena,  it  is  inevitable  that  considerations 
of  space  should  limit  our  statements  to  the  merest  indication  of  the  process  and 
the  briefest  record  of  the  chemical  result.  The  prodigious  disproportion  between 
the  meagerness  of  this  fragment  and  the  mass  of  detail  which  lies  behind  it  in  each 
case  should  be  constantly  before  the  mind  of  the  reader.  The  book  gives  empiri- 
cal knowledge,  the  laboratory  work  and  the  discussion  of  it  furnish  the  only  real 
knowledge.  The  extent  and  nature  of  this  real  knowledge  may  be  shown  in  con- 
nection with  any  action.  As  an  illustration  we  may  point  out  some  of  the  prob- 
lems which  the  heating  of  potassium  chlorate  presents  to  one  who  is  trying  to 
acquire  an  intelligent  acquaintanceship  with  its  chemistry,  and  has  not  previously 
done  the  experiment. 

First,  the  substance  melts.  It  must  be  realized  that  this  is  a  common  occur- 
rence which  does  not  necessarily  imply  any  profound  change,  and  may  be  reversed 
by  cooling.  Later,  the  liquid  appears  to  boil,  and  the  properties  of  a  boiling  sub- 
stance must  be  known.  If  the  observer  has  been  informed  in  advance  that  the 
body  is  homogeneous,  he  must  know  that,  if  it  is  simply  boiling,  it  will  evaporate 
completely  and  leave  nothing  behind,  and  that  the  temperature  required  to  achieve 
this  will  remain  constant  from  the  beginning  to  the  end.  In  order,  therefore,  to 
become  aware  of  the  fact  that  here  decomposition  is  taking  place,  he  must  note  the 
ways  in  which  the  decomposition  of  potassium  chlorate  differs  from  ordinary  boil- 
ing. For  example,  if  it  were  a  case  of  boiling,  he  should  expect  to  find  the  solid 
body  condensing  on  the  sides  of  the  tubes,  and  note  the  fact  that  no  such  conden- 
sation is  observed,  with  the  appropriate  inferences.  He  should  observe  that,  in 
the  later  stages  at  least,  the  escape  of  the  gas  does  not  cease  when  the  flame 
is  removed,  although  this  would  undoubtedly  occur  in  a  case  of  simple  boiling, 
He  must  further  observe  the  changes  in  the  consistency  of  the  material  and  the 
way  in  which  it  finally  becomes  thick  and  may  even  solidify.  Even  the  most  ex- 


OXYGEN 


85 


perienced  investigator  would  have  to  make  many  careful  experiments  before  he 
could  definitely  classify  the  nature  of  the  phenomenon  being  observed.  The  first 
inference  would  probably  be  that  the  phenomenon  was  certainly  not  one  of  mere 
ebullition.  In  some  ways  it  is  like  the  evaporation  of  a  solution,  obtained,  say, 
by  the  melting  of  a  substance  in  its  own  water  of  crystallization.  Yet,  this 
hypothesis  would  not  explain  to  the  thoughtful  observer  even  the  more  obvious 
features  of  the  phenomenon,  for  the  liquid  which  was  acting  as  a  solvent  would 
have  to  be  amazingly  volatile  if  the  absence  of  any  condensation  on  the  walls  of 
the  tube  was  to  be  accounted  for. 

The  illustration  need  not  be  elaborated  further.  These  remarks  are  sufficient 
to  show  that  even  the  simplest  experiment  presents  an  almost  limitless  field  for  the 
discussion  of  important  questions  which  are  more  or  less  common  to  all  chemical 
phenomena.  It  must  be  noted,  also,  that  chemical  change  is  in  itself  not  per- 
ceptible by  the  senses,  and  that  only  physical  properties  and  physical  phenomena 
are  observed  (cf.  pp.  47-49).  The  chemical  facts,  such  as  the  general  nature  of  the 
change,  the  conditions  under  which,  and  the  facility  with  which  it  occurs,  are 
reached  solely  by  inference.  The  above  example  shows  the  ready  and  thorough 
knowledge  of  physics  which  must  be  at  the  command  of  every  individual  effort  to 
study  even  the  simplest  chemical  phenomenon.  It  is  only  when  the  physics  as 
well  as  the  chemistry  of  the  change  have  been  mastered  that  the  "real  knowledge" 
to  which  reference  was  made  above  has  been  gamed. 

Oxygen  from,  Sodium  Peroxide.  —  Oxygen  can  be  obtained 
conveniently  from  sodium  peroxide  Na202 
and  water  H2O  by  means  of  generators  (Fig. 
29)  similar  to  the  acetylene  generators  used 
on  automobiles.  When  the  metal  sodium  is 
burned  in  air,  sodium  peroxide  is  obtained 
as  a  powder.  This  powder,  after  being 
melted,  solidifies  in  compact,  solid  form,  and 
is  sold  as  "oxone."  The  oxone  is  bought  in 
a  small,  sealed  tin  can,  the  ends  of  which 
are  perforated  in  several  places  just  before 
use.  When  the  valve  (B)  is  opened,  so  that 
the  oxygen  escapes,  the  water,  which  fills  the 
generator  almost  to  the  top,  enters  the  can 
(C)  by  the  holes  in  the  bottom  and  interacts 
with  the  oxone.  When  the  valve  is  shut,  the 
gas  continues  to  be  generated  until  it  has 
driven  the  water  down  again  below  the  level 
of  the  bottom  of  the  can:  FIG.  29. 


// 


Skeleton: 
Balanced: 


H2O-»    NaOH  +  O. 
H2O  -»  2NaOH  +  O. 


86  INORGANIC  CHEMISTRY 

This  method  is  convenient  because  it  works  at  room  temperature  and 
can  be  started  and  stopped  at  will.  The  sodium  hydroxide  produced 
is  very  soluble  in  water  and  remains  dissolved.  Note  that  the  name 
sodium  hydroxide  indicates  the  elements  which  compose  it. 

Specific  Properties  of  Two  Kinds,  Physical  and  Chemical. 

—  We  have  learned  that  every  substance  has  its  own  set  of  specific 
properties.  In  describing  a  substance,  it  is  convenient  to  divide 
the  properties  into  two  classes.  The  list  of  substances  with  which 
the  given  substance  can  enter  into  chemical  combination,  for  ex- 
ample, we  place  under  specific  chemical  properties.  Relations  of  the 
substance  to  any  of  the  varieties  of  chemical  change  (p.  21)  belong 
to  this  class. 

On  the  other  hand,  we  do  not  consider  melting  or  boiling  to  be 
chemical  changes,  so  we  place  the  temperatures  at  which  the  substance 
melts  (m.-p.)  and  boils  (b.-p.),  its  color,  etc.  (for  list,  see  p.  40) 
under  specific  physical  properties. 

Properties  of  either  class  may  be  used  for  recognizing  a  substance. 

Specific  Physical  Properties  of  Oxygen.  —  Oxygen  resembles 
air  in  having  neither  color,  taste,  nor  odor.  The  density  of  a  sub- 
stance is,  strictly  speaking,  the  weight  of  1  cubic  centimeter  (1  c.c.). 
In  the  case  of  a  gas,  we  frequently  prefer  to  give  the  weight  of  1000  c.c. 
(1  liter),  at  0°  and  760  mm.  (1  atmosphere)  barometric  pressure. 
For  oxygen  this  weight  is  1.42900  grams  (Morley).  The  corre- 
sponding weight  for  air  is  1.293,  so  that  oxygen  is  slightly  heavier, 
bulk  for  bulk,  than  air  (in  the  ratio  1.105  :  1).  Oxygen  can  be 
liquefied  by  compression,  provided  its  temperature  is  first  reduced 
below  —118°,  which  is  its  critical  temperature  (q.v.).  The  gas  is 
slightly  soluble  in  water,  the  solubility  at  0°  being  4  volumes  of  gas 
in  100  volumes  of  water  (at  20°,  3  : 100). 

The  solubility  of  oxygen  in  water,  although  slight,  is  in  some 
respects  its  most  important  physical  property.  Fish  obtain  oxygen 
for  their  blood  from  that  dissolved  in  the  water.  With  air-breathing 
animals  (like  man),  the  oxygen  could  not  be  taken  into  the  system, 
if  it  did  not  first  dissolve  in  the  moisture  contained  in  the  walls  of 
the  air  sacs  of  the  lungs,  and  then  pass  inwards  in  a  dissolved  state 
to  the  blood. 

Liquid  oxygen,  first  prepared  by  Wroblevski,  has  a  pale-blue 
color.  At  one  atmosphere  pressure,  that  is,  in  an  open  vessel,  it 
boils  at  -182.5°.  Its  density  (weight  of  1  c.c.)  is  1.13,  so  that  it 
is  slightly  denser  than  water.  By  cooling  with  a  jet  of  liquid  hydro- 


OXYGEN 


87 


gen,  Dewar  froze  the  liquid  to  a  snow-like  pale-blue  solid.     A  tube 
of  liquid  oxygen  is  noticeably  attracted  by  a  magnet. 

Six  Specific  Physical  Properties  of  Each  Gas.  —  Although 
every  substance  has  many  physical  properties,  we  shall  mention  only 
those  which  are  used  in  chemical  work,  with  occasionally  the  addi- 
tion of  any  peculiar  or  unexpected  quality.  It  will  aid  the  memory 
to  recall  the  physical  properties  of  a  gas,  if  we  note  that,  as  a  rule, 
only  six  physical  properties  are  mentioned:  (1)  color,  (2)  taste, 
(3)  odor,  (4)  density,  (5)  ease  of  liquefaction,  defined  by  the  critical 
temperature,  (6)  solubility,  usually  in  water  only. 

It  should  be  noted  that,  of  these,  the  first  three  are  never  stated  quantita- 
tively. Taste  and  odor  cannot,  at  present,  be  denned  on  any  absolute  scale. 
Color  could  be  defined  in  terms  of  the  wave-lengths  of  the  light  reflected  and 
transmitted,  and  of  the  relative  intensities  of  each  wave-length,  but  chemists 
seldom  attempt  anything  so  elaborate.  On  the  other  hand,  the  last  three  are 
comparatively  easy  to  measure,  and  are  always  stated  quantitatively.  Yet, 
in  the  case  of  most  substances,  even  familiar  ones,  careful  determinations  have 
never  been  made.  Thus,  for  potassium  chlorate,  the  published  melting-points 
vary  from  324°  to  351°! 

Solubilities  of  Gases  in  Non- Aqueous  Liquids.  —  Iron  and 
steel  are  oiled  to  prevent  rusting.  The  oil,  however,  does  not  pre- 
vent access  of  oxygen  to  the  metal,  for,  contrary  to  the  common 
impression,  gases  are  roughly  ten  times  more  soluble  in  liquids  like 
petroleum  and  alcohol  than  they  are  in  water.  Water  hastens 
rusting,  and  the  oil  excludes  atmospheric  moisture,  since  water  is 
insoluble  in  oils. 

The  following  solubilities,  expressed  as  the  volume  of  the  gas 
(at  760  mm.)  dissolved  by  one  volume  of  the  liquid,  illustrate  this 
point. 


Gas. 

Petroleum. 

Water. 

Alcohol, 
0° 

10° 

20° 

20° 

0° 

0.229 
0.135 
1.31 

0.202 
0.117 
1.17 

0.028 

0.014 
0.901 

0.041 
0.020 
1.797 

0.284 
0.126 
4.329 

Carbon  dioxide      .    . 

Specific  Chemical  Properties  of  Oxygen.  —  The  chemical 
properties  of  pure  oxygen  are  like  those  of  atmospheric  oxygen,  only 
more  pronounced. 


88 


INORGANIC  CHEMISTRY 


Non-metallic  Elements.  —  Sulphur,  when  raised  in  advance  to 
the  temperature  necessary  to  start  the  action,  unites  vigorously  with 
oxygen  (Fig.  30),  giving  out  much  heat  and  pro- 
ducing a  familiar  gas  having  a  pungent  odor, 
sulphur  dioxide  SO2.  This  odor  is  frequently 
spoken  of  as  the  "smell  of  sulphur,"  but  in  reality 
sulphur  itself  has  no  odor,  and  neither  has  oxy- 
gen. The  odor  is  a  property  of  the  compound  of 
the  two.  The  mode  of  experimentation  can  be 
changed  and  the  oxygen  led  into  sulphur  vapor 
through  a  tube.  The  oxygen  then  appears  to 
burn  with  a  bright  flame,  giving  the  same  product 
as  before. 

Phosphorus,  when  set  on  fire,  blazes  in  oxy- 
gen very  vigorously  forming  a  white,  powdery, 
solid  oxide,  phosphorus  pentoxide  PzO$.  Burning 
carbon,  in  the  form  of  charcoal  or  hard  coal, 
glows  brilliantly  and  is  soon  burnt  up.  It  leaves 
an  invisible,  odorless  gas,  carbon  dioxide  CC>2. 
At  high  temperatures,  oxygen  combines  readily 
with  one  or  two  other  non-metals  (e.g.,  silicon, 
boron,  and  arsenic),  and  to  a  small  extent  (1% 
at  1900°)  with  nitrogen.  It  will  not  combine  di- 
rectly with  chlorine,  bromine,  or  iodine,  although 
oxides  of  the  first  and  last  can  be  prepared  by 
using  other  varieties  of  chemical  change.  With  the  six  members  of 
the  helium  family  (q.v.),  of  which  no  compounds  are  known,  and 
with  fluorine,  oxygen  forms  no  compounds. 

S  +  2O  -*  S02. 

2P  +  5O  ->  P205. 

C  +  20  ->  CO2. 

Metallic  Elements.  —  Iron,  as  we  have  seen,  rusts  exceedingly 
slowly  in  air  (diluted  oxygen)  and,  even  when  red-hot,  gives  hammer- 
scale,  the  black  solid  which  is  broken  off  on  the  anvil,  rather  deliber- 
ately. In  pure  oxygen,  a  bundle  of  picture-wire,  if  once  ignited,  will 
burn  with  surprising  brilliancy,  throwing  off  sparkling  globules  of 
the  oxide,  melted  by  the  heat.  This  oxide  is  a  black,  brittle  sub- 
stance, identical  with  hammer-scale,  and  different  from  rust  (ferric 
oxide  Fe2O3).  It  contains,  in  fact,  a  smaller  proportion  of  oxygen 


FIG.  30. 


OXYGEN  89 


than  the  latter,  and  is  called  magnetic  oxide  of  iron  FeaO^  iden- 
tical with  a  well-known  ore  of  iron. 

Skeleton:  Fe  +    O  ->  Fe3O4. 

Balanced:  3  Fe  +  4O  ->  Fe304. 

All  the  familiar  metals,  excepting  gold,  silver,  and  platinum, 
when  heated,  combine  with  oxygen.  Some  combine  more  vigor- 
ously, others  less  vigorously,  than  does  iron.  Oxides  of  the  three 
metals  just  named  can  be  made  by  varieties  of  chemical  change 
other  than  direct  combination. 

Compound  Substances,  if  they  are  composed  largely  or  entirely 
of  elements  which  combine  with  oxygen,  are  able  themselves  to 
interact  with  oxygen.  Usually,  they  produce  a  mixture  of  the  same 
oxides  which  each  element,  separately,  would  give.  Hence,  wood, 
which  is  composed  of  carbon  and  hydrogen  with  some  oxygen,  when 
burnt  in  oxygen,  produces  carbon  dioxide  and  water  (oxide  of  hydro- 
gen) in  the  form  of  vapor.  Again,  carbon  disulphide  burns  readily, 
giving  carbon  dioxide  and  sulphur  dioxide,  just  as  do  carbon  and 
sulphur,  separately.  Ferrous  sulphide  gives,  similarly,  sulphur  diox- 
ide and  magnetic  oxide  of  iron. 

Tests:  A  Test  for  Oxygen.  —  A  test  is  a  property  which,  because 
it  is  easily  recognized  (a  strong  color,  for  example),  or  for  some  other 
sufficient  reason,  is  commonly  employed  in  recognizing  a  substance. 

Oxygen,  as  we  have  seen  (p.  17),  when  pure,  is  recognized  by 
the  fact  that  a  splinter  of  wood,  glowing  at  one  end,  bursts  into  flame 
when  introduced  into  the  gas.  Only  one  other  gas  (see  nitrous 
oxide)  behaves  similarly. 

Determining  Formulse  Again.  —  To  learn  the  exact  nature  of 
interactions  like  those  used  as  illustrations  above,  quantitative  ex- 
periments must  of  course  be  made. 
Thus,  for  example,  a  known  weight  of 
sulphur  is  placed  in  a  porcelain  boat 
(Fig.  31),  which  has  already  been 
weighed.  The  U-shaped  tube  to  the 
right  contains  a  solution  of  potassium 

hydroxide  which  is  capable  of  absorbing  the  resulting  gas.  The 
oxygen  enters  from  the  left.  When  the  sulphur  is  heated,  it  burns 
in  the  oxygen,  and  the  loss  in  weight  which  the  boat  undergoes 
shows  the  amount  of  sulphur  consumed.  The  gain  in  weight  of  the 


90  INORGANIC   CHEMISTRY 

U-tube  shows  the  weight  of  the  compound  produced.  By  subtract- 
ing, we  get  the  quantity  of  oxygen.  The  proportion  of  the  con- 
stituents and  the  steps  in  the  calculation  (p.  72)  are  as  follows: 

PEBCENTAGE  AT.  WT.  FACTOR  -i- 1.561 

Sulphur,         50.05       =       32.06      X     1.561       =     S  X  1.561         S  X  1 
Oxygen,         49.95       =       16.00      X     3.122      =     O  X  3.122        O  X  2 

The  formula  of  the  product  is  therefore  S02  and  the  equation  (p.  73) 
is  S  +  20  -»  SO2. 

Similarly,  phosphoric  anhydride  may  be  shown  to  have  the  for- 
mula P205,  carbon  dioxide  C02,  and  magnetic  oxide  of  iron  Fe304. 

The  results  given  by  the  experiment  described  above  (Fig.  31)  are  usually 
inexact.  The  tendency  to  the  formation  of  sulphur  trioxide,  often  heightened  by 
catalytic  action  (see  below)  of  the  porcelain  of  the  boat,  raises  abnormally  the 
proportion  of  oxygen.  The  principle  of  the  experiment  is  easy  to  understand, 
however. 

In  the  case  of  phosphorus  a  similar  plan  may  be  used.  Instead  of  attempting 
to  receive  the  solid  product  in  a  U-tube,  however,  it  must  be  caught  by  a  plug 
of  glass  wool  in  the  main  tube  of  hard  glass,  and  a  drying  tube  will  be  needed 
at  the  end  to  prevent  admission  of  moisture  from  the  air.  The  increase  in  weight 
of  the  hard  glass  tube  represents  the  oxygen  taken  up.  Care  and  leisurely  per- 
formance are  needed  to  make  the  experiment  successful. 

The  same  method  used  with  sulphur  can  be  employed  for  carbon,  since  the 
carbon  dioxide  is  absorbed  by  the  potassium  hydroxide.  With  carbon,  exact 
results  are  obtained. 

The  data  and  working  in  these  cases  and  in  that  of  iron  are  as  follows: 

PERCENTAGE  AT.  WT.  FACTOR  -r-  0.704 

Phosphorus,  43.66  31.0          X         1.408  P  X  2 

Oxygen,  56.34          =         16.0          X         3.521  O  X  5 

-i-  2.272 

Carbon,  27.27          =         12.0          X         2.272  C  X  1 

Oxygen,  72.72  16.0          X         4.545  O  X  2 

-j-  0.431 

Iron,  72.38          =         55.9          X         1.295  Fe  X  3 

Oxygen,  27.62          =         16.0          X         1.726  O  X  4 

Oxides:  Nomenclature.  —  Substances  containing  one  element 
in  combination  with  oxygen  are  called  oxides.  When  the  same 
element  forms  more  than  one  oxide,  the  names  of  the  oxides  indicate 
the  differing  proportions.  Thus  we  have  barium  oxide  (or  monoxide) 
BaO,  and  barium  peroxide  (or  dioxide)  Ba02,  magnetic  oxide  of 


OXYGEN  91 


iron  Fe3O4,  ferrous  oxide  FeO,  and  ferric  oxide  FezOs.  In  cases  like 
the  last  two  the  terminations  -ous  and  -ic  applied  to  the  metal  cor- 
respond to  the  smaller  and  larger  proportions  of  oxygen,  respectively, 
which  the  metal  is  able  to  hold  in  combination.  The  same  terminations 
are  used  to  distinguish  chlorides,  sulphides  (p.  59),  and  other  com- 
pounds, when  more  than  one  of  each  is  known. 

The  discussion  of  the  formation  and  properties  of  ozone  (q.v.), 
which  is  an  oxide  of  oxygen,  cannot  be  taken  up  until  we  are  in  pos- 
session of  the  means  of  understanding  the  difference  in  density  of 
the  two  substances  (Chap.  XII). 

Combustion.  —  Since  oxygen  is  a  component  of  the  atmos- 
phere, chemical  actions  in  which  it  plays  a  part  are  familiar  in  daily 
life.  Violent  union  with  oxygen  is  called  in  popular  language  com- 
bustion or  burning.  Yet  since  oxygen  is  only  one  of  many  gaseous 
substances  known  to  the  chemist,  and  similar  vigorous  interactions 
with  these  gases  are  common,  the  term  has  no  scientific  significance. 
The  union  of  iron  and  sulphur,  for  example,  gives  out  light  and  heat, 
and  is  quite  similar  in  the  chemical  point  of  view  to  combustion. 

Oxidation.  —  A  number  of  cases  of  the  union  of  an  elementary 
substance  with  oxygen  have  been  referred  to  (pp.  88,  89).  In 
each  case  the  substance  was  heated  strongly  and  the  union  with 
oxygen  was  rapid  (combustion).  Cold  oxygen,  either  pure  or  as  it 
is  found  diluted  in  the  air,  however,  acts  upon  cold  substances  in 
like  manner.  One  difference  is  that  in  the  cold  the  action  is  much 
slower,  and  oxidation  (rusting)  of  the  whole  specimen  may  occupy 
months  or  even  years.  Another  difference  is  that  the  products  of 
slow  and  of  rapid  oxidation  are  not  always  identical  in  composition 
and  properties.  In  the  case  of  iron,  for  example,  burning  gives  us 
the  magnetic  oxide  FeaC^,  while  rusting  in  moist  air  yields  a  hydrated 
ferric  oxide  Fe2O3  +  Aq.*  The  products  differ  in  composition,  but 
are  nevertheless  closely  related. 

This  process  of  slow  oxidation,  although  less  conspicuous  than 
combustion,  is  really  of  greater  interest.  Thus  the  decay  of  wood  is 
simply  a  process  of  oxidation  whereby  the  same  ultimate  products 
(namely,  carbon  dioxide  and  water)  are  formed  as  by  the  more 
rapid,  ordinary  combustion.  Again,  large  volumes  of  pure  water 

*  The  formula  H2O  may  not  be  used  excepting  to  indicate  a  definite  proportion 
of  the  elements  of  water  (18  parts).  Where  the  proportion  varies  according  to 
circumstances,  as  here  and  in  the  case  of  solutions,  the  contraction  Aq.  is  employed 


92  INORGANIC  CHEMISTRY 

are  mixed  with  sewage,  the  object  being,  not  simply  to  dilute  the 
latter,  but  to  introduce  water  containing  oxygen  in  solution.  This 
has  an  oxidizing  power  like  that  of  oxygen  gas  and,  through  the 
agency  of  bacteria,  quickly  renders  dissolved  organic  matters  in- 
nocuous by  converting  them  for  the  most  part  into  carbon  dioxide 
and  water.  Thus  a  few  miles  further  down  the  stream,  the  water 
becomes  as  suitable  for  drinking  as  it  was  before  the  sewage  entered. 

In  our  own  bodies  we  have  likewise  a  familiar  illustration  of  slow 
oxidation.  Avoiding  details,  it  is  sufficient  to  say  that  the  oxygen, 
from  the  air  taken  into  the  lungs,  combines  with  the  haemoglobin 
in  the  red  blood-corpuscles.  When  blood  is  placed  under  the  air 
pump,  it  effervesces,  and  oxygen  is  given  off  (Mayow,  1669).  In 
this  form  of  loose  combination,  it  is  carried  by  the  blood  through- 
out our  tissues  and  there  oxidizes  the  foodstuffs  which  have  been 
absorbed  during  digestion.  The  material  products  are  carbon  diox- 
ide and  water,  of  which  the  former  is  carried  back  to  the  lungs  by 
the  blood-serum,  and  finally  reaches  the  air  during  exhalation.  The 
important  product,  however,  is  the  heat,  given  out  by  the  oxidation, 
which  keeps  the  body  warm.  If  we  cease  to  eat,  we  become  lighter 
and  weaker,  showing  that  a  real  portion  of  our  structure  is  gradually 
being  consumed  by  oxidation. 

The  opposite  of  oxidation,  the  removal  of  oxygen,  is  spoken  of 
in  chemistry  as  reduction.  But  this  term,  as  we  shall  see,  like  oxi- 
dation, has  been  stretched  to  cover  other  kinds  of  chemical  change. 

Uses  of  Oxygen.  —  A  number  of  the  practical  applications  of 
oxygen  have  already  been  mentioned.  For  example,  in  the  fore- 
going section  we  have  referred  to  its  use  in  breathing,  its  role  in 
decay,  which  is  a  beneficent  process  because  it  removes  much  use- 
less matter  which  might  otherwise  cause  disease,  and  its  value  in 
the  disposal  of  sewage.  Power  and  heat  for  commercial  purposes 
are  almost  all  obtained  by  the  burning  of  coal,  in  which  oxygen  from 
the  air  plays  a  large  part.  If  we  had  to  purchase  the  oxygen  as  well 
as  the  coal,  we  should  require  at  least  three  tons  of  oxygen  for  every 
ton  of  coal. 

Oxygen  in  cylinders,  and  from  oxygen  generators,  is  used  to  re- 
store the  supply  in  the  atmosphere  of  submarine  boats,  as  well  as 
for  the  purposes  already  mentioned  (p.  81). 

Substances  Indifferent  to  Oxygen.  —  Finally,  since  the  at- 
mosphere contains  so  large  a  proportion  of  oxygen,  substances  which 


OXYGEN  93 

do  not  oxidize  and,  when  heated,  do  not  burn,  have  many  uses. 
Gold,  silver,  and  platinum  are  of  this  kind  and  are  used  for  ornaments. 
The  last  is  used  for  crucibles  in  which  bodies  are  heated  in  the  labo- 
ratory. Although  iron  burns  in  pure  oxygen,  it  does  not  oxidize 
rapidly  in  the  air  even  when  heated,  and  so  is  u^ed  for  making  ves- 
sels for  cooking  and  in  constructing  fire-proof  buildings. 

Compounds,  already  fully  oxidized,  are  naturally  not  combus- 
tible. Of  this  nature  are  sandstone,  granite,  brick,  porcelain,  glass, 
and  water.  All  these  are,  therefore,  fireproof.  Moreover,  these 
substances  do  not  give  off  oxygen  when  heated  (steam  decomposes 
slightly).  Glass  and  porcelain  thus  neither  lose  nor  gain  in  weight 
when  heated,  and  are  suitable  materials  for  laboratory  apparatus. 

Activity  and  Stability.  —  A  substance  which  enters  in  com- 
bination vigorously,  as  does  oxygen,  is  called  chemically  active. 
Nitrogen,  on  the  other  hand,  is  relatively  inactive.  An  active  ele- 
ment, since  it  combines  eagerly,  naturally  holds  tenaciously  to  the 
matter  with  which  it  has  combined.  An  active  element  implies, 
therefore,  also  one  which  is  in  general  difficult  to  liberate  from  com- 
bination. Its  compounds  are  in  general  relatively  stable.  Thus, 
many  oxides,  and  the  natural  compounds  just  mentioned  (sand- 
stone, granite,  brick,  and  porcelain,  the  last  two  made  from  clay), 
do  not  lose  oxygen  even  at  a  white  heat  and  are  very  stable. 

Means  of  Altering  the  Speed  of  a  Given  Chemical  Action: 
By  Change  of  Temperature.  —  That  the  same  change  may  pro- 
ceed with  very  different  speeds  according  to  conditions  is  a  familiar 
fact.  For  example,  raising  the  temperature  increases  the  rapidity 
of  all  chemical  interactions.  Thus,  cold  iron  combines  with  oxygen 
very  slowly,  giving  rust,  while  white-hot  iron  sheds  quantities  of 
scales  of  an  oxide,  formed  in  the  few  moments  that  it  is  under  the 
blacksmith's  hammer.  White-hot  coal  unites  with  oxygen  in  the 
air  to  form  carbon  dioxide  and  seems  to  disappear  before  our  eyes, 
while  in  the  cellar,  even  in  warm  weather,  we  observe  no  appreciable 
diminution  in  its  amount.  Careful  measurement,  however,  shows 
that,  when  stored  in  the  open  air,  coal  does  lose  from  2  to  5  per 
cent  of  its  heating  value.  When  air  is  largely  excluded,  say,  by 
storage  under  water,  there  is  no  loss.  No  temperature  can  be  found 
at  which  the  interaction  definitely  begins.  We  believe  that  every 
such  change  proceeds  with  some  speed  at  every  temperature.  A 
rough  estimation,  based  on  experiment,  shows  that  on  an  average, 


94  INORGANIC  CHEMISTRY 

other  things  being  equal,  every  rise  in  temperature  of  ten  degrees 
doubles  the  amount  of  material  changed  per  second,  and  conversely. 

If,  on  bringing  two  materials  together,  the  chemist  observes  no 
marks  of  chemical  action,  he  immediately  begins  cautiously  to  heat 
the  mixture.  This  appeal  to  the  magnifying  effect  of  a  rise  in  tem- 
perature is  always  made  as  a  matter  of  course. 

The  common  expressions  used  in  chemistry  in  describing  temper- 
atures, along  with  the  corresponding  readings  of  the  thermometer, 
are  as  follows: 

Incipient  red  heat,  about  525°.     Yellow  heat,  about  1100°. 

Dark  red  heat,  "      700°.     Beginning  white  heat,    "    1300°. 

Bright  red  heat,          "      950°.     White  heat,  "   1500°. 

Rapid  Self-sustaining  Chemical  Action  and  Means  of  Ini- 
tiating it.  —  When  a  piece  of  wood  is  set  on  fire  at  one  end,  the 
heat  produced  by  the  action  itself  raises  the  temperature  of  neigh- 
boring portions  until  their  speed  of  union  becomes  equal  to  that  of 
the  part  originally  lighted.  In  this  way  the  whole  becomes  finally 
inflamed.  When  we  blow  the  blaze  out,  the  great  excess  of  cold  air 
suddenly  lowers  the  temperature  of  the  wood,  and  of  the  gas  rising 
from  it,  and  rapid  union  ceases.  Cold  water,  naturally,  lowers  the 
temperature  even  more  promptly.  Whether  a  given  set  of  materials 
can  maintain  itself  at  a  temperature  proper  to  violent  interaction 
will  depend  on  the  amount  of  heat  developed  by  the  action  itself, 
on  the  one  hand,  and  the  losses  of  heat  by  conduction  and  radiation 
on  the  other.  If  the  latter  are  great,  the  former  must  be  greater. 
Thus  the  union  of  iron  and  oxygen  per  se  gives  heat  enough  to  warm 
the  materials  to  the  burning  temperature  and  leaves  much  over  for 
radiation.  But  iron  in  air,  which  is  four-fifths  nitrogen,  can  receive 
the  oxygen  only  one-fifth  as  fast  at  the  start,  and  even  more  slowly 
as,  later,  the  nitrogen  accumulates  round  it.  And  besides,  all  the 
nitrogen  has  to  be  heated  to,  perhaps,  2000°.  The  task  is  too  great. 
The  union  is  impeded  and  the  iron  is  not  oxidized  fast  enough  to 
generate  the  heat  required  to  maintain  everything  at  this  high 
temperature.  Poor  conductors  of  heat,  like  wood  and  candles,  fare 
better.  Powdered  iron,  with  its  particles  presenting  large  surface  to 
the  air  relatively  to  the  weight  of  material  in  each  particle  to  be 
heated,  burns  well. 

The  initial  supply  of  heat  required  to  start  violent  exothermal 
chemical  actions,  of  which  alone  we  are  here  speaking,  must  not  be 


OXYGEN  95 

contused  with  the  heat  subsequently  developed  a&  the  action  pro- 
ceeds. The  latter  is  usually  much  greater.  Indeed,  the  preliminary 
supply  varies  with  circumstances,  and  may  be  made  as  small  as  we 
choose  by  limiting  the  area  first  heated  and  using  ordinary  precau- 
tions against  radiation  and  convection.  In  practice,  a  single  spark 
from  an  induction  coil  often  takes  the  place  of  more  clumsy  methods 
of  raising  the  temperature.  The  heat  produced  by  the  interaction 
itself,  however,  is  fixed  in  amount,  and  depends  only  on  the  materials 
and  their  quantity. 

Heating  is  not  the  only  means  used  to  give  the  initial  accelera- 
tion to  a  self-sustaining  chemical  change.  The  materials  in  a  match- 
head  are  capable  of  undergoing  a  great  transformation.  Yet,  so 
slowly  does  this  proceed  at  ordinary  temperatures,  that  matches 
may  be  kept  in  efficient  condition  for  years.  Here  a  rather  violent 
vibration  is  employed  to  hasten  the  torpid  action  in  a  small  part  of 
the  material,  and  the  heat  produced  by  the  resulting  action  quickly 
ignites  the  whole. 

The  misleading  term,  kindling  temperature,  is  often  used  in  this  connec- 
tion. It  gives  the  impression  that  there  is  a  definite  temperature  at  which 
combustion  will  start.  But  the  temperature  is  only  one  of  the  conditions  which 
produce  combustion.  Finely  powdered  iron  will  start  burning  at  a  lower  tem- 
perature than  will  an  iron  wire,  because  it  presents  relatively  more  surface  to  the 
gas.  Finely  divided  lead,  known  as  "lead  pyrophorus,"  catches  fire  at  room 
temperature.  Again,  if  the  oxygen  .is  at  less  than  one  atmosphere  pressure,  the 
wire  will  require  to  reach  a  higher  temperature  before  combustion  will  begin. 
Finally,  the  vapor  of  methyl  alcohol  and  air  requires  to  be  raised  above  a  red 
heat  before  combustion  starts,  but  a  pocket  cigar-lighter  sets  fire  to  this  very 
mixture  by  means  of  a  contact  agent  (a  thin  platinum  wire)  without  any  other 
means  of  heating  being  required.  So  that,  the  conditions  under  which  com- 
bustion begins  involve  the  physical  condition  of  the  solid,  the  pressure  of  the 
gas  or  vapor,  the  presence  or  absence  of  a  contact  agent  and  the  nature  of  the 
contact  agent,  as  well  as  the  temperature.  No  definite  kindling  temperature 
can  be  given,  unless  the  other  conditions  are  specified  also.  Kindling  con- 
ditions involve  several  variables,  of  which  the  temperature  is  only  one. 

Spontaneous  Combustion.  —  Sometimes  a  mere  slow  oxida- 
tion develops  into  a  combustion,  which  is  then  known  as  spontaneous 
combustion.  To  understand  this,  we  must  note  the  fact  that  a  given 
weight  of  material,  say,  iron,  in  combining  with  oxygen  to  form  a 
given  oxide,  will  liberate  the  same  total  amount  of  heat  whether  the 
union  proceeds  rapidly  or  slowly.  If  the  action  proceeds  slowly, 
and  the  material  being  oxidized  is  freely  exposed  to  the  air,  the  lat- 


96  INORGANIC  CHEMISTRY 

ter  will  become  heated  and  will  carry  off  the  heat  as  fast  as  it  is 
produced.  Thus,  no  particular  rise  in  temperature  will  occur.  If, 
however,  the  material  is  a  poor  conductor  of  heat,  like  hay  or  rags, 
and  there  is  sufficient  air  for  oxidation,  but  not  enough  to  carry  off 
the  heat,  the  heat  may  accumulate  and  a  temperature  sufficient  to 
start  combustion  may  be  reached.  Such  a  situation  sometimes 
arises  in  hay-stacks.  It  occurs  also  when  rags,  saturated  with  oils 
used  in  making  paints  (linseed  oil  and  turpentine),  are  left  in  a 
heap.  These  oils,  in  "drying,"  combine  with  oxygen  from  the  air 
and  turn  into  a  tough,  resinous  material.  The  rags,  being  poor 
conductors  of  heat,  finally  become  hot  enough  to  burst  into  flame, 
and  serious  conflagrations  often  owe  their  origin  to  causes  such  as 
this.  Oily  rags  should  always  be  disposed  of  by  burning,  or  should 
at  least  be  placed  in  a  closed  can  of  metal.  Fires  in  coal  bunkers 
of  ships  arise  from  the  same  cause,  slow  oxidation  (p.  93),  with  ac- 
cumulation of  the  resulting  heat. 

Other  Means  of  Altering  the  Speed  of  a  Given  Chemical 
Change:  By  Change  in  Concentration;  by  Catalysis;  by  Solu- 
tion. —  Even  when  the  temperature  remains  constant,  there  are 
other  changes  in  the  conditions  (p.  41)  which  may  be  used  for  accel- 
erating or  for  moderating  the  speed  of  chemical  interactions.  The 
most  important  of  these  is,  a  change  in  the  concentration  of  the 
interacting  substances.  Another  is  the  presence  of  a  catalytic  or 
contact  agent.  The  condition  of  solution  might  be  accounted  still 
another. 

The  abatement  in  the  activity  of  the  oxygen  found  in  the  air 
(p.  94),  by  the  nitrogen  which  is  mixed  with  it,  is  a  question  of  con- 
centration. If  the  concentration  of  pure  oxygen  under  atmospheric 
pressure  is  taken  as  unity,  that  of  oxygen  in  air  is  only  about  0.2, 
And  the  speed  of  interaction  of  a  body,  other  things  being  equal,  is 
directly  proportional  to  its  concentration.  This  is  not  an  obscure 
law,  but  merely  common  sense  put  into  definite  language.  The 
opportunity  which  one'  substance  has  for  getting  at  every  part  of 
another  will  be  one  factor  in  determining  the  speed  with  which  the 
resulting  transformation  will  take  place.  And  this  opportunity, 
other  things  being  equal,  depends  on  the  thickness  or  density  with 
which  the  substance  is  scattered  in  the  region  of  action.  In  the 
case  of  a  gas,  this  factor  is  measured  by  its  partial  pressure.  Hence, 
lights  burn  badly  at  great  elevations,  where  the  oxygen  is  very 
tenuous.  On  the  other  hand,  powdered  charcoal,  which  burns  feebly 


OXYGEN  97 

in  common  air,  interacts  so  rapidly  when  ignited  in  liquid  air,  where 
the  oxygen  is  highly  condensed,  that  an  explosion  takes  place.  Again, 
when  oxygen  is  compressed  in  contact  with  barium  oxide  at  700°  it 
combines  to  form  the  dioxide  (p.  82);  when  the  pressure  of  the 
oxygen  in  contact  with  the  latter  is  reduced,  oxygen  is  liberated 
(see  Chemical  equilibrium). 

When,  without  any  change  in  temperature,  an  extra  substance 
increases  the  speed  of  a  chemical  change,  seemingly  by  its  mere 
presence,  without  itself  suffering  any  permanent  change,  we  call 
this  catalytic  (Gk.  Kara,  down,  AVCRS,  the  act  of  loosing)  or  contact 
action.  The  word  was  originally  used  for  cases  of  decomposition. 
The  foreign  body  is  called  the  catalytic  or 
contact  agent,  and  the  process  catalysis. 
The  effect  of  manganese  dioxide  on  the  de- 
composition of  potassium  chlorate  (p.  83) 
is  of  this  nature.  When  some  of  the 
chlorate  is  placed  in  a  flask,  provided  with  a 
two-hole  stopper  and  exit  tube,  and  is 
melted  carefully  so  as  to  avoid  superheat- 
ing, scarcely  any  evolution  of  oxygen  can 
be  perceived  at  this  temperature  (351°). 
If  now  a  pinch  of  pulverized  manganese 
dioxide,  hitherto  held  in  the  closed  tube, 
be  dropped  into  the  molten  mass  by  turning  the  end  of  the  tube 
into  a  vertical  position  (Fig.  32),  the  oxygen  is  given  off  in  torrents 
in  consequence  of  the  enormous  acceleration  of  the  decomposition. 
To  avoid  injury  from  an  explosion,  it  is  advisable  to  wrap  the  flask 
in  a  towel,  before  turning  the  tube.  Yet  the  manganese  dioxide 
may  be  recovered  unchanged  from  the  residue.  Manganese  dioxide, 
of  course,  will  itself  give  oxygen  (p.  82),  but  the  decomposition  is 
hardly  noticeable  at  400°.  Oxone  (p.  85)  always  contains  a  trace  of 
cuprous  hydroxide,  which  hastens  the  action  on  water. 

It  is  found  that  many  actions  owe  what  appears  to  be  their  normal 
speed  to  the  presence  of  a  trace  of  water  vapor.  Thus  many  of  the 
elements  show  no  visible  tendency  to  unite  with  carefully  dried  oxy- 
gen, even  when  they  are  strongly  heated  in  it.  Addition  of  a  trace  of 
moisture,  however,  brings  about  instant  combustion.  So  water  is  to 
be  regarded  as  one  of  the  commonest  contact  agents. 

A  few  cases  of  retardation  of  an  action  by  a  catalytic  agent  are  known.  Thus 
a  little  benzyl  alcohol  or  mannite  added  to  the  solution  will  retard  the  oxidation  of 
sulphites  by  the  air  (Bigelow).  Hence  positive  and  negative  catalysis  both  occur. 


98  INORGANIC  CHEMISTRY 

The  effect  of  solution  in  hastening  a  chemical  change  was  seen 
when  we  examined  the  interaction  of  sodium  chloride  and  silver 
nitrate  (p.  19).  With  the  solutions  the  action  was  seemingly  in- 
stantaneous. If  we  had  attempted  to  bring  it  about  by  rubbing 
the  dry  substances  in  a  mortar,  hours  of  work  would  have  left  much 
of  the  original  bodies  still  unchanged.  Even  heating  would  not 
have  produced  so  prompt  an  effect.  It  is  obvious  that  the  intimate 
access  which  every  part  of  each  solution  gains  to  every  part  of  the 
other  accounts  to  some  extent  for  the  difference  (see  lonization). 
Chemical  actions,  as  will  be  seen  in  the  sequel,  are  very  frequently 
carried  out  in  aqueous  solution  in  order  to  take  advantage  of  the 
favorable  influence  of  this  condition. 

Thermochemistry.*  —  As  we  have  seen  (p.  28)  a  chemical 
change  may  be  accompanied  either  by  a  liberation  or  an  absorption 
of  heat.  Actions,  like  the  oxidations  in  the  present  chapter,  in 
which  heat  is  produced,  are  called  exothermal  actions.  Actions 
which,  like  the  decomposition  of  mercuric  oxide  (p.  17),  or  of 
barium  dioxide  (p.  82),  absorb  heat,  and  proceed  only  so  long  as 
heat  is  furnished,  are  called  endothermal  actions.  Since  the  activ- 
ities, or  affinities  of  two  substances  (say,  two  metals)  may  often  be 
measured  (p.  37)  by  observing  the  amounts  of  heat  liberated  when 
each  combines  with  a  third  substance  (say,  oxygen),  it  will  be  in- 
structive now  to  consider  some  of  the  elementary  facts  of  thermo- 
chemistry. 

The  chemical  interactions  to  be  studied  thermally  are  arranged 
so  that  they  may  be  carried  out  in  some  small  vessel  which  can  be 
placed  inside  another  containing  water.  The  whole  apparatus  is 
called  a  calorimeter.  The  heat  developed  raises  the  temperature 
of  the  water.  Where  gases,  like  oxygen,  are  concerned,  a  closed 
bulb  of  platinum  forms  the  inner  vessel.  The  quantity  of  heat 
capable  of  raising  one  gram  of  water  one  degree  in  temperature,  at 
15°,  is  called  a  calorie.  So  that  250  grams  of  water  raised  1°  would 
represent  250  calories,  and  20  grams  of  water  raised  5°  would  repre- 
sent 100  calories. 

While  in  physics  the  unit  of  quantity  is  the  gram,  in  chemistry 
the  unit  which  we  select  is  naturally  that  represented  by  the  formula 
of  the  substance.  Thus,  the  heat  of  combustion  of  carbon  means 
the  heat  produced  by  combining  twelve  grams  of  carbon  (charcoal) 

*  This  section  may  be  omitted  at  this  point,  and  can  be  taken  up  very 
appropriately  in  connection  with  fuels,  under  carbon. 


OXYGEN  99 

with  thirty-two  grams  of  oxygen,  and  is  sufficient  to  raise  97,600 
grams  of  water  one  degree.  This  is  expressed  as  follows: 

C  +  2O  -»  C02  +  97,600  cal. 

In  other  words,  the  combustion  of  less  than  half  an  ounce  of  carbon 
will  raise  one  kilogram  (over  two  pounds)  of  water  from  0°  almost 
to  the  boiling-point. 

It  is  always  found  that  the  same  quantities  of  any  given  chemical 
substances,  sustaining  the  same  chemical  change  under  the  same 
conditions,  produce  or  absorb,  according  as  the  action  is  exothermal 
or  endothermal,  amounts  of  heat  which  are  equal. 

The  rate  at  which  a  given  chemical  action  is  allowed  to  take  place 
has  no  influence  on  the  total  amount  of  heat  consumed  or  produced. 
It  may  not  at  first  sight  appear  obvious  that  rusting  evolves  heat, 
but  a  delicate  thermometer  will  show  that  a  heap  of  rusting  nails  is 
somewhat  higher  in  temperature  than  surrounding  bodies.  As  we 
have  seen  (p.  96),  poor  conductors,  like  oily  rags  and  ill-dried  hay, 
show  a  tendency  to  spontaneous  combustion  owing  to  accumulation 
of  the  slowly  developing  heat  of  oxidation;  and  the  warmth  of  our 
own  bodies  is  due  to  the  same  cause. 

In  accordance  with  invariable  experience  expressed  in  the  law  of 
the  conservation  of  energy,  when  an  action  is  chemically  capable  of 
reversal,  the  contribution  of  the  same  amount  of  heat  which  it  de- 
velops will  exactly  suffice  to  drive  the  chemical  change  in  the  opposite 
direction.  The  heat  contributed  is  simply  used  to  restore  the  amount 
of  chemical  energy  proper  to  the  original  system.  Thus,  the  union 
of  one  chemical  unit  weight  each  of  mercury  and  oxygen  (p.  80) 
produces  30,600  cal. : 

Hg  +  O  ±5  HgO  +  30,600  cal., 

and  the  decomposition  of  one  formula-weight  of  mercuric  oxide  (p.  17) 
demands  the  same  amount  of  heat  in  order  that  free  mercury  and 
oxygen,  with  their  appropriate  proportions  of  internal  energy,  may 
be  recovered. 

In  practice  it  is  found  that  all  chemical  changes  are  not  capable  of 
reversal  by  the  use  of  the  sources  of  heat  available  in  the  laboratory. 
A  quantity  of  heat,  equivalent  to  that  produced  by  any  chemical 
action  on  a  small  scale,  is  very  easily  provided,  but  something  more 
appears  to  be  necessary.  The  heat  provided  must  be  of  a  certain 
temperature,  otherwise  it  is  quite  ineffective.  For  example,  the  heat 


100  INORGANIC  CHEMISTRY 

produced  by  the  union  of  calcium  and  oxygen  is  within  the  limits  of 
ready  measurement, 

Ca  +  0  ->  CaO  +  131,000  caL, 

and  the  supply  of  this  amount  (or  even  of  unlimited  amounts)  of  heat 
to  calcium  oxide  (quicklime)  is  easily  achieved.  Yet  this  method  is 
quite  ineffective  to  produce  decomposition  (p.  23)  of  the  product. 
Apparently  we  have  not  sufficiently  high  temperatures  for  the  pur- 
pose at  our  command  (see  Factors  of  energy). 

It  may  be  noted  in  this  connection  that  the  temperatures  required  to  produce 
reasonably  rapid  decomposition  vary  within  a  vide  range.  Some  substances 
can  be  kept  only  below  0°,  and  Hecompose  when  allowed  to  become  warm.  Others, 
like  the  oxides  of  gold  and  platinum,  require  a  little  heating  (p.  122).  Many, 
like  quicklime,  are  not  broken  up  even  at  the  temperature  of  the  electric  arc. 
When  the  energy  is  applied  in  the  form  of  electricity  (p.  28),  instead  of  heat,  the 
range  is  incomparably  more  easily  within  the  reach  of  the  means  ordinarily  at  our 
disposal.  There  is  no  substance,  provided  it  is  of  such  a  nature  as  to  be  affected 
by  the  electric  current  at  all,  which  cannot  be  decomposed  by  a  current  with  an 
E.M.F.  of  10  volts  or  less,  while  currents  of  110  volts  and  over  are  commonly 
accessible.  It  is  partly  on  this  account  that  electrical  processes  have  become  so 
common  hi  industrial  chemistry. 

One  of  the  most  important  principles  of  thermochemistry  is  the 
law  of  constant  heat  summation.  If  a  system  of  substances  can  be 
transformed  into  another  system  of  substances  by  different  stages  or 
by  more  than  one  route,  then  the  algebraic  sum  of  the  heats  absorbed 
or  produced  in  the  various  routes  is  the  same.  Thus,  barium  oxide 
might  be  formed  either  directly  from  the  proper  proportions  of  the 
constituents,  or  indirectly  by  preparing  the  dioxide  (p.  82),  and  then 
driving  out  half  of  the  oxygen  contained  in  the  latter.  The  quan- 
tities of  heat  involved  in  these  two  routes  are  as  follows: 

Direct:  Ba  +  O  -»  BaO  +  124,400  cal. 

(  Ba  +  20  -*  Ba02  +  141,600  cal.  (1) 

(  Ba02  ^  BaO  +  O  -  17,200  cal.  (2) 

Ba  +  O  -^  BaO  +  124,400  cal.  (3) 

When  equations  (1)  and  (2)  are  added  algebraically,  canceling  terms 
such  as  BaO2  and  O,  which  are  common  to  both  sides  of  the  final 
equation,  the  fact  that  the  indirect  route  is  exactly  equivalent  to 
the  direct  one  is  at  once  apparent,  and  the  balance,  in  favor  of  heat 
produced,  is  124,400  calories  as  before.  If  in  such  cases  the  sum 


OXYGEN  101 

of  the  heats  were  not  the  same,  it  would  follow  that  by  using  different 
plans  of  procedure  we  could  prepare  .different  specimens  of  the  same 
substance  containing  different  proportions  of  chemical  energy. 
This,  however,  we  have  never  been  able, to  do. 

The  quantities  of  heat  liberated  in  two  chemical  changes  are  often 
measures  of  the  relative  amounts  of  available  chemical  energy  in  the 
systems  before  the  change,  and,  therefore,  often  furnish  a  measure  of 
the  relative  chemical  activities  of  the  two  sets  of  substances.  The 
comparison  may  safely  be  made  in  certain  cases  when  the  conditions 
under  which  the  two  actions  take  place  are  precisely  alike.  Formerly 
it  was  supposed  that  the  heat  liberated  was  always  proportional  to 
the  chemical  activity  of  the  substances,  but  we  have  already  shown 
cause  (pp.  35,  37)  why  this  general  statement  cannot  be  true. 

It  should  be  noted  that  production  or  absorption  of  heat  is  not, 
in  itself,  an  evidence  of  chemical  action.  Physical  changes  are 
likewise  accompanied  by  the  same  phenomena.  Thus,  the  evapora- 
tion of  water  absorbs  heat,  and  condensation  of  a  vapor  and  the 
crystallization  of  a  supercooled  liquid  liberate  heat. 

Exercises.* — 1.  Enumerate  other  instances,  already  encount- 
ered, of  the  use  of  the  terminations  ous  and  ic  to  distinguish  different 
degrees  of  oxidation.  For  what  other  purposes  have  the  same 
terminations  been  used? 

2.  What  difference  in  composition  between  potassium  chloride 
and  chlorate  are  the  terminations  designed  to  indicate?    Applying 
the  same  idea,  how  would  ferrous  sulphate  (q.v.)  differ  from  ferrous 
sulphide,  and  cupric  sulphate  from  cupric  sulphide? 

3.  Define  and  illustrate:  density  of  a  gas  (p.  86,  and  see  p.  Ill), 
density  of  a  solid  or  liquid  (pp.  3,  86). 

4.  Enumerate  the  classes  of  facts  given  under  the  heads  of, 
Physical  Properties,  and  Chemical  Properties  of  oxygen,  respectively. 

5.  Construct  the  equations  for  the  combustion  of  phosphorus, 
carbon,  and  iron  in  oxygen  (pp.  90,  91). 

6.  What  weight  of  oxygen  may  be  obtained  by  heating  30  g.  of 
barium  dioxide  (p.  82)?     In  solving  this  problem,  pay  minute  atten- 
tion to  the  rules  (p.  75). 

7.  When  1  g.  of  sodium  burns  in  oxygen,  it  produces  1.7  g.  of  the 
oxide.     What  is  the  formula  of  the  latter,  and  the  equation  (p.  72)? 

8.  How  should  you  show  that,  in  the  making  of  oxygen  from  a 
mixture  of  potassium  chlorate  and  manganese  dioxide,  the  latter 

*  See  footnote  to  p.  26. 


102  INORGANIC  CHEMISTRY 

remains  unchanged?     WTricn  properties  (p.  40)  are  you  employing 
for  this  purpose? 

9.  The  substances,  like  phosphorus  and  sulphur,  which  burn 
rapidly  in  ordinal  y  ox-ygenf  combine  very,  very  slowly  (even  when 
heated)  with  oxygen  which  has  been  freed  from  moisture  by  careful 
drying.     How  is  this  effect  of  water  to  be  classified? 

10.  Discuss  the  union  of  iron  and  sulphur  (p.  16)  and  the  decom- 
position of  mercuric  oxide  (p.  17)  in  their  relation  to  the  explana- 
tions on  pp.  94r-95. 

11.  How  many  calories  are  required  to  raise  500  g.  of  a  substance 
of  specific  heat  0.5  from  15°  to  37°  (p.  98)? 

12.  The  combustion  of  1  g.  of  sulphur  to  sulphur  dioxide  de- 
velops 2220  calories.     What  is  the  heat  of  combustion  of  sulphur 
(p.  99)? 

13.  Outline  briefly  the  proof  that  thermochemical  data  are  not 
accurate  measures  of  chemical  activity  (p.  35). 

14.  What  percentage  by  weight  of  free  oxygen  is  obtained  by 
heating:    (a)  mercuric  oxide,   (6)  potassium  nitrate,   (c)  potassium 
chlorate?    At  $1.50  (6/8),  $0.15  (8d),  and  $0.15  (8d)  per  kilogram, 
respectively,  which  is  the  cheapest  source  of  oxygen? 

15.  Using  the  data  on  p.  86,  calculate  the  weight  of  oxygen  dis- 
solved by  100  c.c.  (=  100  g.)  of  water  at  20°. 

16.  Why  does  a  forced  draft  make  a  fire  burn  more  rapidly? 

17.  Why  does  a  naked  flame  sometimes  cause  an  explosion  in  a 
mine,  when  the  air  of  the  mine  is  filled  with  coal  dust? 

18.  When  iron  wire  burns  in  oxygen,  the  oxide  of  iron  falls  in 
melted  globules,  and  no  smoke  is  produced,  while  when  phosphorus 
is  burned,  all  the  oxide  goes  off  as  smoke,  and  nothing  remains  in  the 
spoon.     Explain  this  difference. 


CHAPTER  VI 
THE  MEASUREMENT   OF   QUANTITY  IN   GASES 

WE  have  spoken  of  measuring  the  proportion  by  weight  of  the 
oxygen  used  in  several  chemical  changes,  but  in  our  illustrations  we 
have  never  weighed  the  gas  itself.  We  have  always  (e.g.,  p.  90) 
obtained  its  quantity  by  subtracting  the  weights  of  solid  or  liquid 
bodies.  In  practice  this  method  often  serves  the  purpose. 

Our  preference  for  weighing  as  a  means  of  ascertaining  quantity 
of  matter  is  largely  due  to  the  fact  that  the  weight  is  independent  of 
the  physical  or  even  chemical  condition  of  the  substance.  Yet,  with 
proper  precautions,  we  may  learn  the  quantity  of  matter  by  means  of 
any  other  attributes  which  are  proportional  to  it.  Now  the  volume 
is  such  an  attribute.  In  determining  the  quantity  of  a  liquid,  where 
rapidity  with  no  extreme  accuracy  is  desired,  the  volume  is  frequently 
measured.  In  the  case  of  gases  the  error  made  in  measuring  the 
volume  is  less,  as  a  rule,  than  in  measuring  the  weight. 

The  Variable  Concentration  of  Gases.  —  A  little  experience 
with  gases  soon  shows  us  that  measurement  of  volume  alone  does 
not  necessarily  give  any  definite  idea  of  the  quantity  of  matter 
present.  The  denseness  with  which  the  gaseous  matter  is  packed 
(the  concentration  of  the  gas)  in  the  vessel  must  somehow  be  defined, 
as  well  as  its  volume,  in  order  that  there  may  be  specification  of  the 
quantity  of  matter. 

Gases  vary  markedly  in  chemical  activity  with  changes  in  their  concentration 
(cf.  p.  94),  and  thus  the  consideration  of  this  condition  (p.  41)  is  continually 
forced  upon  the  chemist.  Solids  and  liquids  do  not  alter  their  denseness  of  pack- 
ing (concentration)  very  noticeably  even  when  compressed  severely  or  changed  in 
temperature.  So  that  concentration  need  not  be  considered  in  the  case  of  pure 
bodies  in  the  solid  or  liquid  forms.  Such  substances  can  be  scattered  through  a 
variable  space  by  solution  in  some  solvent,  however,  and  then  their  degree  of 
packing  or  concentration  becomes  an  important  factor  in  their  chemical  behavior 
also. 

The  principle  used  for  estimating  the  concentration  of  a  gas  may 
be  illustrated  by  means  of  the  arrangement  in  Fig.  33.  Except  tha.t 

103 


104 


INORGANIC  CHEMISTRY 


a  little  gas  (any  gas  will  do)  remains  shut  off  by  the  mercury  in  the 
left  limb  of  the  tube,  the  whole  apparatus  has  been  evacuated.  The 
reservoir  can  be  turned  upward,  and  thus  larger  amounts  of  mercury 
may  be  introduced  into  the  tube. 

Now  the  portion  of  the  mercury  below  the  dotted  line  is  essen- 
tially a  balance,  that  is  to  say,  it  will  move  in  one  direction  or  the 
other  if  the  stresses  on  either  side  change. 
At  present  these  stresses  must  be  equal. 
On  the  right  pan  of  the  balance,  so  to 
speak,  the  stress  is  represented  by  the 
weight  of  the  column  of  mercury  above 
the  dotted  line.  As  there  is  nothing  in  the 
tube  above  this  mercury,  the  weight  of  the 
latter  is  all  that  this  side  of  the  balance 
sustains.  On  the  left  pan  of  the  balance 
there  must  be  an  equal  stress,  and  this 
stress  can  be  caused  only  by  the  gas  con- 
fined in  the  shorter  limb.  The  nature  of 
a  gas  suggests  that  this  stress  must  be  ex- 
ercised on  the  walls  of  the  tube  also,  al- 
though they  naturally  do  not  exhibit  its 
effects.  This  stress  we  call  the  pressure  of 
the  gas. 

The  height  of  the  surface  of  the  mercury 
on  the  right  above  that  on  the  left  having 
been  measured,  more  mercury  may  now  be 
added  from  the  reservoir,  and  the  difference 
in  the  two  levels  again  noted.  The  gas  can- 
not have  diminished  in  amount,  yet  it  now 
occupies  a  smaller  space,  and  is,  therefore, 
packed  more  closely  —  its  concentration  is  greater  than  before.  If, 
for  example,  the  difference  in  level  is  now  twice  as  great,  it  will  be 
found  that  the  concentration  of  the  gas  is  also  twice  as  great  (its 
volume  having  become  half  of  the  original  volume).  Whatever 
amount  of  mercury  is  added,  we  shall  always  find  that  the  con- 
centration of  the  gas  is  proportional  to  the  height  of  the  mercury. 
But  this  in  turn  is  proportional  to  the  weight  of  metal.  The  weight 
of  mercury  on  one  side  must,  therefore,  be  equal  to  the  stress  or 
pressure  or  tension  of  the  gas  on  the  other  side  which  balances  it. 
Hence,  the  concentrations  of  a  sample  of  any  gas  are  proportional 
to  the  corresponding  pressures  it  exercises.  We  determine,  therefore, 


FIG.  33. 


THE  MEASUREMENT  OF  QUANTITY  IN  GASES 


105 


the  denseness  with  which  any  sample  of  gas  is  packed  by  measuring 
its  pressure. 

Method  of  Allowing  for  Varying  Concentration  in  Measur- 
ing Quantity  in  Gases.  —  The  principle  just  stated  is  applied  to 
the  measurement  of  the  quantity  of  matter  in  a  sample  of  gas  by  per- 
mitting the  concentration  of  the  sample  to  alter  until  it  is  equal  to 
that  of  the  atmosphere  at  the  moment.  Then  we  read  off  the  volume 
now  occupied,  and  simultaneously  we  ascertain  the 
pressure  by  observing  that  of  the  atmosphere. 
Each  of  these  two  operations  is  facilitated  by  a 
special  arrangement  of  apparatus. 

A  gas  to  be  measured  is  always  confined  so 
that  some  liquid  constitutes  one  part  of  the  barrier 
to  its  escape.  The  very  simplest  form  of  the  ap- 
paratus is  shown  in  Fig.  34.  To  render  the  con- 
centration (and  pressure)  of  the  gas  equal  to  that 
of  the  atmosphere,  the  cylinder  containing  the  gas 
is  lowered  (or  raised)  until  the  levels  of  the  liquid 
inside  and  outside  are  the  same.  When  the  sys- 
tem is  in  this  condition  the  stress  of  the  gas  on  the 
inner  surface  must  be  equal  to  that  of  the  atmos- 
phere on  the  outer  one,  otherwise  movement  of 
the  liquid  would  occur.  The  volume  of  the  gas 
is  then  read  directly  from  the  graduation  on  the 
cylinder.  Often  the  cylinder,  or  other  vessel,  is 
closed  with  a  ground-glass  plate,  placed  quickly  in 
erect  position,  and  weighed.  The  weight  of  water 
which  is  then  required  to  fill  it  to  the  brim  gives 
more  exactly  the  volume  occupied  by  the  gas  (1  g. 
water  =  1  c.c.).  When  special  modes  of  admitting 
or  handling  the  gas  have  to  be  provided  for,  the  apparatus  may  be 
more  complex.  But  the  principle  of  adjustment  is  always  the  same. 
In  exact  work,  mercury  is  employed  to  confine  the  gas.  Water 
serves  the  purpose  of  rough  work  with  gases  which,  like  oxygen,  are 
but  slightly  soluble  in  it.  When  water  is  used,  the  volume  is  too 
great  by  the  space  occupied  by  the  vapor  of  the  water  which  is  mixed 
with  the  gas  (see  Mixed  gases,  p.  Ill),  and  a  correction  must  always 
be  made  on  this  account. 

The  pressure  or  tension  of  the  atmosphere  at  any  moment  is 
measured  by  means  of  a  simplified  form  of  the  apparatus  in  Fig.  33. 


FIG.  34. 


106 


INORGANIC   CHEMISTRY 


The  reservoir  is  omitted,  but  the  space  above  the  mercury  on  the 
right  is  still  a  vacuum.  The  atmospheric  air  being  the  gas  whose 
concentration  is  to  be  measured  through  its  pressure,  the  short  limb 
is  left  open.  The  resulting  apparatus  (Fig.  35)  performs  its  functions 
in  the  same  way  as  does  the  more  complex  one.  The  only  difference 
is  that  mercury  is  automatically  added  to  or  with- 
drawn from  the  right  side  by  the  motion  of  the 
metal  resulting  from  changes  in  the  pressure  of  the 
air.  The  reading  of  the  vertical  height  between  the 
lower  and  upper  surfaces  of  the  mercury  gives  a 
number  which  is  proportional  to  the  weight  of  mer- 
cury on  the  right  side  of  the  balance  and,  therefore, 
to  the  (equal)  stress  of  the  atmosphere  on  the  left. 
This  is  called  the  barometric  reading  (uncorrected), 
after  the  name  of  the  instrument. 

To  make  different  readings,  taken  when  the 
mercury  is  at  different  temperatures,  strictly  pro- 
portional to  the  weight  of  the  metal,  the  observed 
height  is  always  reduced  to  that  which  would  have 
been  shown  by  the  same  weight  of  mercury  at  0°  in 
the  same  apparatus.  A  thermometer,  and  a  table 
of  temperatures  with  the  corresponding  corrections 
to  be  subtracted  from  the  uncorrected  reading  (C, 
Fig.  35),  must  be  used. 

Knowing  now  the  volume  occupied  by  the  sample 
of  gas  when  its  concentration  is  equal  to  that  of  the 
atmosphere  and  the  barometric  reading,  which  is 
proportional  to  this  concentration,  the  measurement  of  the  amount 
of  matter  in  the  sample  has  become  definite  so  far  as  concerns  the 
variability  of  concentration  with  change  in  pressure. 

Boyle9 s  Law.  —  The  recorded  results  of  measurements  made  as 
above  at  different  times,  when  the  atmospheric  pressures  are  different, 
are  still  unsatisfactory  because  the  data  for  samples  of  the  same  kind 
of  gas  differ  in  the  value  of  the  pressure  as  well  as  in  that  of  the 
volume.  To  make  the  results  easily  comparable  in  respect  to  the 
amount  of  matter  they  represent,  one  further  step  is  needed.  All 
the  data  are  recalculated  so  as  to  show  the  volume  each  sample 
would  have  occupied  if  the  pressure  had  been  equal  to  the  weight  of 
760  mm.  of  mercury,  which  is  the  average  height  of  the  barometer  at 
the  sea  level  in  45°  of  latitude. 


Fia.  35. 


THE   MEASUREMENT  OF  QUANTITY   IN  GASES  107 

We  have  seen  that  the  concentration  of  a  given  quantity  of  a  gas 
is  proportional  to  its  pressure  (p.  104).  But,  volume  occupied  is  the 
inverse  of  concentration.  Thus,  the  same  rule  may  be  stated  in  the 
form :  The  volumes  occupied  by  a  sample  of  any  gas  are  inversely  pro- 
portional to  the  pressure  at  each  volume.  The  fact  was  discovered 
by  Boyle  (1660)  who  stated  it  in  this  way.  In  still  other  words,  the 
product  of  the  several  pressures  and  corresponding  volumes  of  a 
sample  of  gas  is  a  constant. 

A  numerical  illustration  will  show  the  mode  of  applying  this  rule. 
We  measure  200  c.c.  of  a  gas  at  atmospheric  pressure,  and  the  barom- 
eter reads  742  mm.  The  question  is:  What  would  be  the  volume 
of  this  amount  of  gas  at  760  mm.  barometric  pressure?  It  will  be 
200  X  y|§  c.c.  =  volume  at  760  mm.  It  is  unnecessary  to  use  any 
formula,  but  absolutely  essential  to  ask:  Is  the  new  pressure  greater 
or  less  than  the  old?  Here  it  is  greater.  Hence,  according  to  the 
law,  the  new  volume  will  be  less,  so  that  the  fraction  must  be  arranged 
with  the  smaller  number  in  the  numerator. 

Boyle's  law  may  be  illustrated  by  using  a  long  tube  bent  like  a  barometer  (Fig. 
35)  but  having  the  short  limb  closed  and  the  long  one  open.  Strips  of  paper  mark 
the  levels  of  the  mercury,  which  are  at  first  alike  on  both  sides,  and  register  the 
volume  of  the  air  in  the  short  limb  at  a  pressure  of  one  atmosphere.  The  reading 
of  the  barometer  at  the  time,  say  750  mm.,  is  noted.  Then  mercury  is  added 
through  a  funnel  inserted  in  the  long  limb,  until  the  level  in  this  limb  is  750  mm. 
above  that  in  the  other.  A  stick  cut  to  750  mm.  length,  and  held  beside  the  tube, 
will  conveniently  show  when  this  has  been  accomplished.  The  pressure  in  the 
open  limb  being  now  two  atmospheres,  the  volume  of  the  air  will  be  found  to  have 
diminished  to  one-halt  its  former  value. 

For  pressures  lower  than  one  atmosphere,  a  different  arrangement  must  be 
used.  A  graduated  tube,  closed  at  one  end,  is  partly  filled  with  mercury  and  in- 
verted in  a  tall,  narrow  cylinder  full  of  the  same  metal.  The  tube  is  then  clamped 
in  any  position,  such  that  the  mercury  level  in  the  tube  is  above  that  in  the  cylin- 
der. The  reading  of  the  barometer  is  next  noted.  The  volume  occupied  by  the 
air  in  the  tube  is  then  read,  and  the  difference  in  height  of  the  two  mercury  sur- 
faces is  measured  by  means  of  a  graduated  rule.  Subtracting  this  height  from 
that  of  the  barometer  gives  the  pressure  of  the  air  in  the  tube.  The  position  of 
the  tube  is  then  altered,  and  the  same  measurements  repeated,  as  often  as  is 
wished.  The  product  of  each  volume  and  the  corresponding  pressure  will  be  a 
constant  number.  The  law  is  expressed  mathematically  by  letting  pi  and  Vi 
represent  the  initial  pressure  and  volume,  p2  and  vz  the  new  pressure  and  volume, 
and  so  forth.  Then  p&i  =  pzvz  =  constant  for  that  particular  specimen  of  gas. 
For  a  given  sample  of  gas,  any  one  of  the  four  values  may  be  calculated  if  the  other 
three  are  known, 


108 


INORGANIC   CHEMISTRY 


Boyle's  law  states  the  facts  with  sufficient  accuracy  for  all  ordi- 
nary purposes.  But  in  reality  no  two  gases  behave  precisely  alike  in 
respect  to  change  in  concentration  when  the  pressure  is  altered  (see 
Chap.  IX).  The  same  gas  does  not  even  behave  in  precisely  the 
same  way  at  high,  intermediate,  and  low  pressures.  The  ideal  gas, 
which  should  behave  uniformly,  we  call  the  perfect  gas.  With  most 
gases,  at  low  pressures  concentration  increases  more,  and  at  very 
high  pressures  much  less  than  the  rule  indicates  (see  Kinetic  theory). 

The  Relation  of  the  Volume  of  a  Gas  to  Temperature.  — 

In  the  foregoing  we  have  assumed  that  there  were  no  temperature 
effects.  But,  as  a  matter  of  fact,  a  rise  in  temperature  will  at  once 
produce  an  increase,  and  a  fall  in  temperature  a  decrease 
in  the  pressure  of  an  enclosed  sample  of  a  gas.  Hence  a 
record  of  the  pressure  alone  will  fail  to  indicate  the  con- 
centration definitely,  and  volume  and  pressure  together 
will  still  leave  the  amount  of  material  unspecified.  The 
temperature  must,  therefore,  be  given  as  well. 

Our  descriptions  of  different  samples  of  gas,  at  dif- 
ferent temperatures,  having  thus  become  once  more 
incomparable,  we  apply  the  same  kind  of  remedy  as 
before.  We  calculate  the  volume  which  each  specimen 
of  gas  would  occupy  at  0°. 

The  rule  for  this  calculation  may  be  demonstrated  in 
a  rough  way  as  follows :  The  large,  graduated  bulb  (Fig. 
36)  is  surrounded  by  a  vessel  which  can  subsequently 
be  filled  with  ice  water  or  with  water  of  any  tempera- 
ture up  to  100°.  About  one-half  of  the  bulb  is  occu- 
pied by  the  gas.  The  mercury  which  fills  the  rest  is  connected  with 
a  reservoir,  so  that  the  levels  of  the  metal  can  be  made  alike,  and  the 
pressure  of  the  gas  be  maintained  constantly  the  same  as  that  of  the 
atmosphere.  When,  now,  the  volume  occupied  by  the  gas  at  0°  is 
read,  and  warmer  water  is  introduced,  we  find  that  the  volume  gains 
-$}?  of  its  value  at  0°  for  every  degree  through  which  its  temperature 
rises.  If  it  is  cooled  below  0°,  it  loses  ^  of  its  volume  at  0°  for 
every  degree  through  which  the  temperature  is  lowered.  Observa- 
tion gives  practically  the  same  value  for  all  gases. 

The  following  graphic  method  will  put  these  facts  in  a  clearer 
light.  In  Fig.  37  we  have  on  the  left  a  thermometer  scale  divided 
into  degrees  Centigrade.  The  middle  line  represents  the  volumes  of 
a  given  sample  of  gas  which  correspond  with  the  successive  tem- 


FJG.  36. 


THE   MEASUREMENT  OF  QUANTITY  IN  GASES 


109 


peratures.  If  we,  for  convenience,  take  a  volume  of  273  c.c.  of  a  gas 
at  0°  and  warm  this  through  1°,  then  at  1°  its  volume,  having  gained 
2^  °f  ^s  original  value,  becomes  274  c.c.  At  2°  it  has  gained 
another  ^^  of  the  volume  it  had  at  0°  and  becomes  therefore  275 
c.c.,-etc.  If  cooled  below  0°  it  loses  ?fa  of  the  volume,  becoming 
272  c.c.,  etc. 


Temp. 
Cent. 


7' 
6' 
5< 
4< 
3( 
2< 

r 

o< 
-r 


Vol.  of 
273  c.c. 


-  275  c.c. 

-  274  c.c. 

-  273  c.c. 

-  272  c.c. 

-  271  c.c. 


Temp. 
Abs. 


I-  100°  373  c.c. 

99°       M  372  c.c. 

[-   98°       \-  371  c.c. 

I  I 


373° 
372° 
371° 


275° 

274° 
273° 
272° 
271° 


I 


-271° 

-272° 
-273° 


FIG.  37. 


I 


2° 
1° 
0° 


If  the  gas  be  heated  to  100°  it  gains  %%%  of  its  volume  at  0°  and 
becomes  373  c.c.  We  might  infer  that  if  it  were  cooled  through  273° 
from  0°  it  would  lose  |^f  of  its  volume,  in  other  words,  it  would  dis- 
appear. This  temperature  has  not  yet  been  reached,  and  in  any 
case  all  gases  would  presumably  liquefy  before  reaching  it.  Our 
statements  apply  to  ordinary  temperatures  only,  and  within  them 
the  law  holds  with  considerable  strictness.  Now  the  series  of  num- 
bers on  the  middle  line  (Fig.  37)  are  all  273  units  larger  than  the 
temperatures  Centigrade  in  the  line  to  the  left  of  the  figure.  If  we 
alter  the  graduation  of  the  Centigrade  thermometer  by  adding  273 
to  each  temperature,  we  secure  the  scale  on  the  right,  which  exhibits 
degrees  of  the  very  same  length  as  before,  but  has  the  additional 
advantage  that  the  numbers  expressing  temperature  are  the  same  as 


110  INORGANIC  CHEMISTRY 

those  expressing  volume.  This  is  what  we  call  the  absolute  scale  of 
temperature.  In  this  artificial  case  we  started  with  273  c.c.  at  0°  C. 
(273°  Abs.),  and  the  absolute  temperature  is  here  always  numerically 
equal  to  the  volume.  If  a  different  volume  had  been  taken  at  0°, 
then  the  volume  assumed  by  the  gas  at  each  temperature  would 
have  borne  a  constant  ratio  to  the  volumes  recorded.  Hence,  the 
volumes  assumed  by  a  sample  of  gas  at  different  temperatures,  the 
pressure  remaining  constant,  are  in  the  same  proportion  as  the  corre- 
sponding absolute  temperatures.  This  is  the  modern  way  of  stating 
a  fact  which  was  first  discovered  by  Charles  of  Paris  (1787).  Obvi- 
ously, if  the  volume  remains  constant,  then  the  pressure  will  be  pro- 
portional to  the  absolute  temperature.  The  bare  fact  underlying 
our  statement  of  the  law  is  that  all  gases  suffer  equal  increments  (or 
decrements)  in  volume  (or  pressure)  for  equal  changes  in  temperature. 

The  discovery  of  this  fact  is  generally  attributed  to  Dalton,  who  first  published 
an  investigation  of  the  subject,  embodying  this  result,  in  1801,  or  to  Gay-Lussac, 
who  made  a  more  complete  investigation  in  1802.  Inasmuch,  however,  as  in 
Chemistry  we  have  another  important  principle  known  as  Gay-Lussac' s  law,  and 
several  which  are  connected  with  the  name  of  Dalton,  it  is  on  the  whole  fortunate 
that  we  are  justified  in  attributing  this  discovery  to  Charles. 

The  application  of  this  law  may  best  be  illustrated  by  an  example. 
We  obtain  200  c.c.  of  a  gas  at  17°  and  wish  to  know  what  volume  it 
would  occupy  at  0°.  To  answer  the  question,  we  convert  the  Centi- 
grade temperatures  to  the  absolute  scale  by  adding,  algebraically, 
273  to  each.  Thus  200  X  fH  =  the  volume  at  0°  required.  No 
formula  is  needed.  We  simply  ask  whether  the  new  temperature  is 
higher  or  lower  than  the  old  one.  Here  it  is  lower.  The  new  volume 
will  therefore  be  smaller  than  the  old  one.  So  we  take  care  to  place* 
the  smaller  number  in  the  numerator. 

The  law  may  be  put  in  mathematical  form.  It  is  preferable,  however,  that 
beginners  should  use  the  method  employed  in  the  illustration  given  in  the  pre- 
ceding paragraph.  Simple  mathematical  expressions,  like  the  one  representing 
this  law,  are  not  made  to  save  us  the  trouble  of  remembering  the  law  itself, 
and  it  would  be  unfortunate  if  their  use  led  us  to  forget  it. 

The  behaviors  of  gases  in  respect  to  changes  of  temperature  and 
pressure  are  perfectly  independent  of  one  another,  so  that  the  above 
laws  may  be  applied  to  any  example,  either  in  succession,  using  the 
answer  for  the  first  calculation  in  making  the  second,  or  simultane- 
ously. Thus  200  c.c.  of  gas  at  742  mm.  pressure  and  17°  become 
200  X  JH  X  Ht  =  183.8  c.c.  at  0°  and  760  mm. 


THE   MEASUREMENT  OF  QUANTITY   IN  GASES  111 

Mixed  Gases.  —  Two  gases  at  the  same  temperature,  provided 
they  do  not  interact  chemically,  do  not  interfere  with  each  other's 
pressures  when  mixed.  Thus,  if  they  are  forced  into  the  same  volume, 
the  pressure  of  the  mixture  is  equal  to  the  sum  of  those  of  the  com- 
ponents (Dalton's  law,  1807).  The  gases  are  therefore  still  thought 
of  individually,  and  the  share  which  each  gas  has  in  the  total  pressure 
is  called  its  partial  pressure.  This,  like  any  other  gaseous  pressure, 
is  proportional  to  the  concentration  of  the  particular  gas  in  the 
mixture. 

For  example,  a  gas  measured  over  water  contains  water  vapor. 
The  partial  pressure  of  this,  called  the  aqueous  tension,  which  is 
definite  for  each  temperature,  must  be  subtracted  from  the  total 
pressure.  The  remainder  is  the  partial  pressure  of  the  gas  being 
measured,  and  this  remainder  is  used  as  the  pressure  of  this  gas  in 
any  calculation.  Thus,  in  a  gas  measured  over  water  at  22°,  the 
total  pressure  includes  19.7  mm.  pressure  of  water  vapor  (see  Appen- 
dix IV).  Hence  150  c.c.  of  gas  over  water  at  22°  and  750  mm.  is 
the  same  amount  as  150  c.c.  of  the  same  gas  in  dry  condition  at  22° 
and  730.3  mm.,  there  being  simply  150  c.c.  of  water  vapor  at  19.7 
mm.  mixed  with  it.  To  obtain  the  volume  of  diy  gas  at  0°  and 

273      730  3 

760  mm.,  we  have  the  expression  150  X  — -  X  • 

dvO  i  DU 

Densities  of  Gases.  —  The  application  of  the  laws  of  Boyle  and 
Charles  enables  us  to  express  the  quantities  of  matter  in  samples  of 
gases  in  a  definite  and  readily  comparable  manner.  In 
describing  chemical  changes,  however,  it  is  continually 
necessary  to  express  quantities  of  gases  by  weight.  The 
relation  between  volume  and  weight  for  each  kind  of  gas 
must,  therefore,  be  ascertained.  If  we  know,  for  ex- 
ample, the  weight  of  one  liter  of  each  gas  at  0°  and  760 
mm.  pressure,  conversion  to  other  weights,  volumes, 
temperatures,  and  pressures  can  be  made.  The  one- 
thousandth  part  of  this  value,  the  weight  of  1  c.c.,  is 
called  the  density  of  the  gas.  Often,  however,  the  rel- 
ative weights  of  equal  volumes,  with  that  of  air  or 
hydrogen  as  unity,  receive  this  name. 

For  most  chemical  purposes  a  high  degree  of  accuracy  is  not  re- 
quired. The  most  direct  method  is  to  employ  a  light  flask  of  125-150 
c.c.  capacity,  provided  with  a  rubber  stopper  and  stopcock  (Fig.  38). 
By  means  of  an  air-pump  the  contents  of  the  flask  are  removed,  and 


112  INORGANIC  CHEMISTRY 

it  is  weighed.  This  gives  the  weight  of  the  empty  vessel.  The  gas, 
whose  density  is  to  be  ascertained,  is  then  admitted,  and  care  is 
taken  that  it  finally  fills  the  flask  at  the  pressure  of  the  atmosphere. 
The  flask  is  closed  and  weighed  again.  The  increase  represents  the 
weight  of  the  gas.  At  the  same  time  the  temperature  and  baro- 
metric pressure  are  read.  The  volume  is  determined  by  displacing 
the  gas  once  more  from  the  flask,  filling  with  water,  and  weighing 
again.  The  difference  in  weight  between  the  empty  flask  and  the 
flask  full  of  water,  in  grams,  represents  the  volume  of  the  content  of 
the  flask  in  cubic  centimeters.  This  volume  is  reduced  to  0°  and 
760  mm.  by  the  rules  discussed  above,  and  we  have  then  a  volume 
of  the  gas  and  the  corresponding  weight. 

To  illustrate,  let  us  suppose  that  the  volume  of  the  flask  is  200 
c.c.  and  that  it  is  filled  with  oxygen  at  17°  and  742  mm.  The  weight 
of  the  gas  is  found  to  be  0.26  g.  We  ascertained  (p.  110)  by  calcula- 
tion that  at  0°  and  760  mm.  this  volume  would  be  183.8  c.c.  The 
weight  of  a  liter  is  given  by  the  proportion  183.8  :  0.26  : :  1000  :  x. 
Here  x  =  1.415  g.  When  the  operation  is  performed  carefully,  and 
the  weighing  carried  to  the  nearest  milligram  instead  of  the  nearest 
centigram,  a  result  more  nearly  approaching  the  exact  one  (1.429) 
may  easily  be  reached. 

To  get  the  density  of  oxygen  referred  to  hydrogen  as  unity,  we 
must  divide  the  answer  by  the  weight  of  a  liter  of  hydrogen  (0.08987 
g.).  In  the  above  case  the  quotient  is  15.74.  The  accepted  value  is 
15.90.  The  density  referred  to  air  as  unity  is  similarly  obtained  by 
dividing  by  1.293,  the  weight  of  a  liter  of  air  at  0°  and  760  mm. 
pressure. 

If  a  suitable  pump  is  not  available,  the  flask,  in  this  case  provided  with  two 
openings,  is  weighed  without  preliminary  exhaustion.  This  gives  the  weight  of 
the  vessel  plus  that  of  the  air  it  contains.  A  continuous  stream  of  the  gas  is  then 
passed  into  the  flask  until  the  air  has  been  completely  displaced.  The  vessel  is 
then  closed  and  another  weighing  made.  Finally  the  gas  is  displaced  by  water, 
and  a  third  weighing  taken.  The  temperature  and  barometric  pressure  are  noted 
as  usual.  The  last  weighing  gives  the  volume  as  before.  Knowing  that  one  liter 
of  air  weighs  1.293  g.  at  0°  and  760  mm.,  we  may  calculate  readily  the  weight  of 
the  air  which  ,the  flask  contained  at  the  observed  temperature  and  pressure. 
When  this  is  subtracted  from  the  number  obtained  in  the  first  weighing,  we  have 
the  weight  of  the  empty  flask.  Subtracting  this  in  turn  from  a  second  weighing, 
we  have  the  weight  of  the  gas.  We  obtain  thus  the  weight  of  a  known  volume 
of  the  gas  at  a  kftown.  temperature  and  pressure,  and  finish  the  calculation  as 
before, 


THE  MEASUREMENT  OF  QUANTITY  IN  GASES  113 

The  values  of  the  densities  of  gases  are  of  great  significance  in  the 
chemical  point  of  view.  A  number  of  them  are  given  in  connection 
with  the  discussion  of  molar  weights  (q.v.). 

Vapor  Densities  of  Liquids  and  Solids.  —  The  densities  of 
vapors  are  as  important  to  the  chemist  as  those  of  gases  and,  solids 
and  liquids  being  more  numerous,  are  even  more  frequently  measured. 

A  modification  of  the  flask  just  described  is  used.  A  tempera- 
ture sufficiently  high  to  vaporize  the  substance  must  be  employed. 
.The  volume  is  reduced  by  rule  to  0°  and  760  mm.,  and  the  density, 
in  this  case  known  as  the  vapor  density  (wt.  of  1  c.c.)  is  calculated 
as  before.  The  reduction  to  0°  and  760  mm.  pressure  by  rule  gives, 
of  course,  a  fictitious  result.  The  vapor  would  condense  to  the 
liquid  form  before  0°  was  reached,  if  the  cooling  were  actually  car- 
ried out.  But  the  value  for  the  density  as  it  would  be  at  0°  and  760 
mm.  has  to  be  calculated  to  facilitate  comparison  with  the  corre- 
sponding values  for  other  substances.  The  results  have  no  physical 
significance,  but  are  highly  important  to  the  chemist. 

Exercises.  —  The  foregoing  cannot  be  understood  unless  some 
problems  involving  the  laws  of  gases  are  actually  worked. 

1.  Reduce  189  c.c.  of  gas  at  15°  and  750  mm.  to  0°  and  760  mm. 

2.  Reduce  110  c.c.  of  gas  at  -5°  and  741  mm.  to  0°  and  760  mm. 

3.  Convert  500  c.c.  of  gas  at  25°  and  700  mm.  to  18°  and  745  mm. 

4.  Reduce  250  c.c.  of  gas,  standing  over  water,  at  21°  and  765 
mm.,  to  the  dry  condition  and  to  0°  and  760  mm. 

5.  An  evacuated  bulb  weighs  13.312  g.     After  being  filled  with 
the  vapor  of  carbon  tetrachloride  at  100°,  it  weighs  13.951  g.     Filled 
with  water  it  weighs  141.3  g.     The  barometric  reading  (corr.)  is 
755  mm.     What  is  the  vapor  density? 

6.  The  density  of  a  substance  referred  to  air  is  3.2.     What  is 
the  density  referred  to  hydrogen?     What  will  be  the  volume  occu- 
pied by  10  g.  of  the  substance  at  20°  and  752  mm.? 


CHAPTER  VII 
HYDROGEN 

HYDROGEN,  although  discovered  by  Paracelsus  in  the  sixteenth 
century,  was  confused  with  other  combustible  gases,  and  its  inde- 
pendent nature  was  first  established  by  Cavendish  in  1766.  Some- 
what later  (1781),  the  latter  showed  that  hydrogen  when  it  burned 
in  the  air  gave  water  vapor,  of  which  he  condensed  a  large  quantity 
to  the  liquid  form.  Taken  in  conjunction  with  Lavoisier's  proof 
that  oxygen  was  the  active  substance  in  the  air  (1777),  this  fact 
showed  that  water  was  a  compound  of  hydrogen  (Gk.  Z8<ap,  water; 
,  to  produce)  and  oxygen,  and  not  a  simple  substance. 


Occurrence.  —  The  free  element  is  found,  mixed  with  varying 
proportions  of  other  gases,  in  exhalations  from  volcanoes,  in  pockets 
found  in  certain  layers  of  the  rock-salt  deposits,  and  in  some  mete- 
orites. The  air  contains  not  more  than  one  part  of  it  in  1,500,000. 
Its  lines  are  very  prominent  in  the  spectra  of  the  sun  and  of  most 
stars. 

In  combination,  it  constitutes  about  11  per  cent  of  water.  It  is 
an  essential  constituent  of  all  acids.  It  is  contained  also,  in  combi- 
nation with  carbon,  in  the  components  of  natural  gas,  petroleum, 
and  all  animal  and  vegetable  bodies. 

We  have  seen  (p.  24)  that,  using  the  physical  unit  of  weight,  the  element 
hydrogen  stands  ninth  in  order  of  plentifulness.  But  the  chemical  work  that 
elements  can  do  should  rather  be  reckoned  by  the  relative  numbers  of  atomic 
weights  which  are  available.  When  Clarke's  numbers  are  recalculated  to  this 
basis,  and  the  number  of  chemical  unit  weights  of  oxygen  is  called  100,  hydrogen 
assumes  a  position  more  in  harmony  with  its  importance: 

Oxygen  .  .  100.00  Aluminium.  .  8.63  Iron  ....  2.37 
Hydrogen  .  31.13  Sodium  ...  3.25  Calcium  ...  2.53 
Silicon  .  .  29.52  Magnesium  .  2.78  Potassium  .  .  1.91 

Preparation  by  the  Action  of  Metals  on  Cold  Water.  —  To 

liberate  hydrogen  from  water,  it  is  necessary  to  use  some  element 
with  which  the  oxygen  of  the  water  will  combine  even  more  eagerly 

114 


HYDROGEN  115 

than  with  hydrogen,  and  to  offer  this  element  in  exchange  for  the 
hydrogen. 

The  most  active  metals,  such  as  potassium  K,  sodium  Na, 
or  calcium  Ca,  displace  hydrogen  rapidly  from  cold  water.  Potas- 
sium and  sodium  are  lighter  than  water,  and  float  on  the  surface. 

In  the  case  of  the  former,  so  much  heat  is  liberated         

that  the  hydrogen  catches  fire,  and  with  neither  metal 
is  the  experiment  safe  in  the  hands  of  a  novice.  Cal- 
cium sinks  to  the  bottom,  and  acts  rapidly,  but  not 
violently,  so  that  the  gas  is  easily  collected  (Fig.  39). 
The  pieces  of  these  metals,  of  course,  act  upon  only 
a  small  part  of  the  water  in  the  vessel.  In  each  case 
the  metal  displaces  one-half  only  of  the  hydrogen  in  p 
that  part  of  the  water  upon  which  it  acts.  The  other  ! 


products  are  the  hydroxides  of  potassium,  sodium,          t,,IG  39> 
and  calcium,  respectively.     The  two  former  dissolve, 
leaving  a  clear  liquid  when  the  metal  is  all  gone,  but  may  be  re- 
covered as  white  solids  by  evaporation.     The  calcium  hydroxide 
(slaked  lime)  is  dissolved  only  in  part,  and  much  of  it  may  be  seen 
after  the  action,  suspended  in  the  water. 

An  alloy  of  lead  with  sodium  (35  per  cent),  sold  under  the  name 
of  hydrone,  affords  a  convenient  substitute  for  sodium  in  the  fore- 
going actions. 

The  Making  of  Equations.  —  To  make  an  equation  we  must 
have  the  results  of  quantitative  measurements.  These  furnish  us 
with  the  composition  of  each  substance  concerned.  The  composi- 
tion, expressed  in  multiples  of  the  atomic  weights,  is  recorded  in  the 
formula  of  the  substance.  If  we  are  in  possession  of  the  necessary 
formulae,  we  can  write  the  equation. 

For  example,  the  composition  of  water  is:  hydrogen  2  X  1.008, 
oxygen  16.  In  symbols,  this  is  2H  and  O,  and  the  formula  is  there- 
fore H2O.  The  composition  of  potassium  hydroxide  is:  potassium 
39.1,  oxygen  16,  hydrogen  1.008,  and  the  formula  therefore  KOH. 
In  calcium  hydroxide  the  proportions  are:  calcium  40.07,  oxygen 
2  X  16,  hydrogen  2  X  1.008  and  the  formula  Ca(OH)2. 

To  make  the  equation,  we  first  write  down  the  formulae  of  the 
substances  used  and  produced: 

K  +  H2O-»KOH  +  H. 
Na  +  H2O->NaOH  +  H. 
Skeleton:  Ca  +  H2O  -^  Ca(OH)2  +  H. 


116  INORGANIC  CHEMISTRY 

Next  we  must  balance  the  equation,  if  necessary.  That  is  we 
must  adjust  it  so  that  there  are  equal  numbers  of  atomic  weights 
(or  atoms)  of  each  element  on  both  sides  of  the  equation.  This  is 
necessary  only  in  the  third  equation,  and  is  done  because,  according 
to  the  law  of  conservation  of  mass,  there  must  be  the  same  quantity 
of  each  element  after  the  reaction  as  there  was  before  it.  On  exam- 
ining the  third  equation,  we  note  there  is  2O,  in  the  (OH)2,  on  the 
right  side  and  only  O  on  the  left.  We  therefore  place  a  2  in  front  of 
the  H20,  for  we  cannot  get  the  additional  oxygen  excepting  by 
using  more  water: 

Balanced:  Ca  +  2H2O  -+  Ca(OH)2  +  2H. 

The  number  of  atomic  weights  of  hydrogen  is  made  equal  by  using 

2H  on  the  right  side. 

The  reader  must  practice  the  making  of  equations,  until  he  can 

do  it  quickly.     The  text  contains  many  equations,  but  more  usually 

only  the  data  required  for  mak- 
ing them  (the  formulae  of  the 
substances)  are  given. 

Hydrogen  from  Metals  and 
Water  at  a  High  Temperature. 

—  With  steam  at  a  red  heat, 
metals  like  iron,  zinc,  and  mag- 
nesium, also  red-hot,  interact  vig- 
orously. The  steam,  generated 
4(X  in  a  flask,  enters  at  one  end  of 

the   tube   containing  the  metal 

(Fig  40),  and  the  hydrogen  passes  off  at  the  other.  Since,  at  a  red 
heat,  all  hydroxides,  except  those  of  potassium  and  sodium,  are  de- 
composed into  an  oxide  of  the  metal  and  water,  as,  for  example, 
Mg(OH)2  —  >  MgO  +  H2O,  the  oxides  are  formed  in  this  case: 

Mg  +  H2O-*MgO  +  2H. 
Iron  gives  the  magnetic  oxide 


Making  Equations,  Again.  —  The  skeleton  equation  for  the 
action  of  iron  on  steam  is: 

Skeleton:  Fe  +  H2O  ->  Fe304  +  H. 

We  are  not  permitted  to  alter  these  formulas  themselves,  but  we 
may  put  coefficients  in  front  of  any  of  them  to  make  the  number  of 


HYDROGEN  117 

atomic  weights  alike  on  both  sides.  The  rule  (p.  73)  is  to  pick  out 
the  largest  formula  and  reason  back  from  that.  Here,  this  is  Fe3O4. 
To  get  Fes,  we  must  start  with  3Fe,  and  to  get  64,  we  must  start 
with4H2O: 

Balanced:  3Fe  +  4H20  ->  Fe304  +  8H. 

Acids.  —  In  making  hydrogen  in  the  laboratory,  the  acids  are 
used  almost  exclusively.  The  common  acids  are  hydrochloric  acid 
HCl,Aq  and  sulphuric  acid  H2S04,Aq.  The  usual  forms  are  mix- 
tures containing  water,  the  variable  amount  of  the  latter  being  indi- 
cated by  the  symbol  Aq.*  Hydrochloric  acid  is  a  solution  of  a  gas, 
hydrogen  chloride.  The  "pure  concentrated"  hydrochloric  acid  used 
in  laboratories  contains  nearly  as  much  of  the  gas  (37  per  cent  by 
weight)  as  the  water  can  dissolve.  When  heated,  it  readily  gives 
up  part  of  the  gas,  and  the  effervescence  attending  this  must  not  be 
mistaken  for  evidence  of  chemical  action.  The  "commercial"  acid 
contains  impurities  and  is  also  less  concentrated.  The  "concen- 
trated," pure  sulphuric  acid  is  an  oily  liquid  containing  practically  no 
water.  The  "commercial"  sulphuric  acid  contains  6  to  7  per  cent  of 
water,  besides  impurities.  Acetic  acid  HCO2CH3,Aq  is  a  solution  of 
a  liquid  in  water. 

All  the  dilute  acids  contain  70  to  80  per  cent  of  water.  The  water, 
as  a  rule,  takes  no  part  in  the  chemical  changes  in  which  the  acids 
are  concerned,  and  is  therefore  omitted  from  the  equations. 

The  name  acid  is  restricted  to  one  class  of  substances  having  cer- 
tain definite  characteristics.  Hydrogen  is  the  one  essential  con- 
stituent of  all  acids.  Their  aqueous  solutions  have  a  sour  taste  and 
change  the  color  of  litmus,  a  vegetable  coloring  matter,  from  blue  to 
red.  When  free  from  water  they  do  not  conduct  electricity.  We 
shall  note  presently  two  other  properties  which  acids  show  when  dis- 
solved in  water,  namely,  that  they  conduct,  and  are  decomposed  by 
the  electric  current,  and  that  their  hydrogen  (or  one  unit  weight  of 
it  in  the  case  of  acetic  acid)  is  displaced  by  certain  metals. 

Radicals.  —  In  describing  the  chemical  behavior  of  acids,  we 
speak  of  the  hydrogen  as  the  positive  radical,  because  in  electrolysis 
(p.  28)  it  is  attracted  to  the  negative  pole,  and  of  the  material 
combined  with  the  hydrogen  as  the  negative  radical  because  it  is 
attracted  to  the  positive  pole  (see  p.  120).  Thus  the  negative  radi- 

*  See  footnote  on  p.  91. 


118  INORGANIC  CHEMISTRY 

cals  in  the  above  acids  are  Cl,  S04,  and  CO2CH3,  respectively.  The 
first  (Cl)  is  a  simple  radical,  the  others  are  compound  radicals.  In 
many  interactions  the  compound  radicals  move  as  units  from  one 
state  of  combination  to  another. 

Preparation  by  Displacement  from  Diluted  Acids.  —  Every 
one  of  the  metals  which  displace  (p.  115)  hydrogen  from  water  will 
also  displace  it  from  dilute  acids.  The  acids  must  be  diluted  with 
water,  unless,  like  hydrochloric  acid,  they  are  already  dissolved  in 
water.  The  action  is  much  more  vigorous  than  that  on  water,  so 
that  the  most  active  metals  are  not  employed.  Metals  like  zinc, 
iron,  and  aluminium  serve  the  purpose.  The  metal  combines  with 
the  negative  radical,  and  so  liberates  the  hydrogen,  which  escapes 
in  bubbles.  Evaporation  of  the  clear  liquid,  when  the  metal  has  all 
disappeared,  gives  in  dry  form  the  compound  of  the  metal  with  the 
negative  radical.  Thus,  with  zinc  and  dilute  sulphuric  acid,  zinc 
sulphate  ZnS04  is  produced : 

Skeleton:  Zn  +  H2SO4  ->  ZnSO4  +  H. 

Balanced:  Zn  +  H2SO4  -> ZhSO4  +  2H. 

With  aluminium  and  hydrochloric  acid,  the  product  is  aluminium 
chloride  A1C13: 

Skeleton:  Al  +  HC1    ->A1C13  +  H. 

Balanced:  Al  +  3HC1  -»  A1C13  +  3H. 

The  water  undergoes  no  change  during  the  action,  although  its  pres- 
ence is  essential.  It  is  simply  a  part  of  the  apparatus.  Any  acid 
may  be  used,  although  with  many  the  action  goes  on  veiy  slowly. 

For  preparing  small  amounts  of  hydrogen,  the  apparatus  (Fig. 
41)  is  such  that  additional  acid  may  be  added  through  the  thistle, 
or  safety  tube.  This  avoids  opening  the  flask  and  admitting  air. 
The  gas  may  be  caught  like  oxygen  over  water  or,  being  lighter  than 
air,  may  be  collected  by  downward  displacement  of  the  latter  (Fig. 
42a).  Heavy  gases  are  collected  by  upward  displacement  of  air 
(Fig.  426). 

With  Kipp's  apparatus  (Fig.  43),  the  gas  may  be  made  on  a 
large  scale  and  its  delivery  can  be  regulated.  When  the  stream  of 
gas  is  shut  off  by  the  stopcock,  the  pressure  of  the  gas,  as  it  continues 
to  be  generated,  drives  the  acid  away  from  the  metal  and  up  into 
the  globe  above,  so  that  the  action  ceases.  Yet  the  action  is  ready 


HYDROGEN 


119 


to  begin  again  the  moment  any  portion  of  the  stored  gas  is  drawn 
off  for  use. 


FIG.  41. 


FIG.  42a. 


FIG.  426. 


Silver,  gold,  and  platinum,  which  do  not  combine  with  free  oxy- 
gen, and  even  copper  and  mercury,  which  do,  are  all  unable  to  form 
oxides  and  to  liberate  hydrogen  when  heated  in 
steam.  When  treated  with  dilute  acids,  none  of 
these  metals  are  able  to  displace  and  liberate  the 
hydrogen  (see  Order  of  activity  of  the  metals,  p. 
129). 

Contact  of  the  zinc  or  iron  with  an  inactive 
metal,  like  platinum  or  copper,  forms  an  electrical 
couple  (q.v.)  and  hastens  the  interaction.  For  the 
same  reason,  commercial  zinc,  which  contains  traces 
of  other  metals,  gives  a  steady  evolution  of  hydro- 
gen, while  extremely  pure  zinc  is  almost  inactive. 

When  water  is  not  used  along  with  the  acid,  the 
latter  is  either  inactive  or  undergoes  a  different 
sort  of  chemical  change.  Thus,  dry,  gaseous,  or 
liquefied  hydrogen,  chloride  hardly  interacts  at  all 
with  zinc.  Pure,  concentrated  sulphuric  acid,  on 
the  other  hand,  although  almost  unaffected  by  zinc 
in  the  cold,  is  violently  decomposed  when  heated. 
The  action,  however,  is  not  a  simple  displacement 
of  the  hydrogen.  The  oxygen  is  removed  from  a  part  of  the  acid, 
and  water  and  hydrogen  sulphide  are  formed: 

4Zn  +  5H2S04  -»  4ZnS04  +  4H2O  +  H2S. 

Preparation  of  Hydrogen  by  Electrolysis.  —  When  we  desire 
to  liberate  hydrogen  by  the  direct  application  of  energy  (p.  28),  we 
find  that  electrical  energy  serves  the  purpose  best. 


FIG.  43. 


120 


INORGANIC  CHEMISTRY 


The  common  compounds  of  hydrogen,  like  hydrogen  chloride  and  water,  are 
not  easily  decomposed  by  heat,  and  in  most  cases,  at  best,  a  mixture  of  gases 
would  be  obtained.  The  difficulty  in  separating  the  resulting  gases  makes  the  use 
of  this  form  of  energy  unsuitable.  On  account  of  its  ability,  not  only  to  liberate 
the  constituents  from  combination,  but  also  to  deliver  the  positive  and  the  nega- 
tive parts  of  the  compound  in  separate  places,  electricity  alone  is  available. 

If  we  dissolve  any  acid  in  water,  and  immerse  the  wires  from  a 
battery  in  the  solution,  bubbles  of  hydrogen  begin  to  appear  on  the 
negative  wire  (the  cathode)  and  rise  to  the  sur- 
face. All  the  other  constituents,  whatever  they 
may  be,  form  the  negative  radical  (see  above) 
and  are  attracted  to  the  positive  wire  (the 
anode)  and  some  of  them  are  set  free  at  its  sur- 
face. An  apparatus  devised  by  Hofmann  (Fig. 
44)  enables  us  to  secure  the  hydrogen,  which 
ascends  on  the  left  and  accumulates  at  the  top 
of  the  tube,  displacing  the  solution.  The  other 
products,  if  gaseous,  occupy  a  separate  tube  on 
the  right  side.  The  solution  displaced  by  the 
gases  is  forced  down  and  mounts  into  the  bulb 
behind.  The  current  of  electricity  flows  from 
one  wire  to  the  other  through  the  liquid  in  the 
cross-tube.  In  a  typical  case,  the  production  of 
hydrogen  ceases  when  the  acid  is  all  decomposed. 
The  water  alone  is  an  almost  complete  noncon- 
ductor, so  that  the  flow  of  the  electricity  prac- 
tically ceases  at  the  same  time.  If  the  operation 
does  not  come  to  rest  in  this  way,  its  contin- 
uance is  due  to  the  regeneration  of  conducting 
substances  by  the  interaction  (with  the  water) 
of  the  materials  of  the  radical  liberated  at  the 
positive  electrode. 

When  hydrochloric  acid  is  used,  we  have  a 
close  approximation  to  the  typical  case.     The  equation  is: 

HC1-+H  (neg.  wire)  +  Cl  (pos.  wire), 

and  the  chlorine,  a  soluble  gas,  remains  dissolved  in  the  water  near 
one  pole.     When  sulphuric  acid  is  employed,  the  equation  is: 


FIG.  44. 


H2SO4— >2H  (neg.  wire)  +  SO4  (pos.  wire), 


(1) 


HYDROGEN  121 

and  the  864  interacts  with  the  water  (see  Discharging  potentials), 
thus: 

S04  +  H2O  -»  H2S04  +  0.  (2) 

Hence  oxygen  comes  off,  and  the  substance  regenerated  is  here  sul- 
phuric acid  itself.  The  final  results  are,  therefore,  the  liberation  of 
hydrogen  and  of  oxygen  and  the  localization  of  the  regenerated  acid 
round  the  positive  electrode. 

Decomposition  of  a  compound  by  the  use  of  electrical  energy  is 
called  electrolysis  (Gk.  kXturpov,  amber;  AvW,  to  loosen). 

It  is  worth  noting  that  the  acids  and  water,  taken  separately,  are 
all  nonconductors.  The  fact  that  the  mixture  does  conduct,  con- 
comitantly  with  the  decomposition  of  the  acid,  is  therefore  highly 
suggestive.  Solution  in  such  cases  must  be  something  more  than  a 
mere  physical  change  of  state  of  aggregation  (see  lonization). 

It  is  commonly  asserted  that  water  is  decomposed  by  a  current  of  electricity. 
This  is  true  in  the  sense  in  which  we  might  say  that  a  man  can  carry  off  a  hill. 
He  may  eventually  remove  it,  if  you  give  him  time.  The  action  of  electricity  upon 
the  purest  water  is  exceedingly  slow,  on  account  of  the  very  minute  conductivity 
for  electricity  which  it  possesses.  Common  distilled  water  owes  its  appreciable 
capacity  for  conducting  chiefly  to  traces  of  an  acid,  namely,  carbonic  acid,  which 
it  contains.  Even  when  the  water  is  saturated  with  carbonic  acid,  however, 
dilute  sulphuric  acid  has  a  conductivity  of  the  order  of  a  thousand  times  greater. 
For  our  present  purpose,  therefore,  water  is  declared  to  be  a  nonconductor.  Yet, 
as  we  shall  see,  the  conductivity  of  pure  water,  small  as  it  is,  has  to  be  taken  into 
consideration  in  certain  cases  (see  Hydrolysis  and  Electromotive  chemistry). 

When  dilute  sulphuric  acid  is  employed,  the  sulphuric  acid  is  regenerated 
by  interaction  with  the  water,  so  that,  in  the  end,  only  the  water  is  decomposed. 
But  the  illustration  is  needlessly  complicated  by  the  chemical  interaction  of  the 
SO4  with  the  water.  The  source  of  the  hydrogen  and  the  oxygen  liberated  is  also 
rendered  obscure  by  the  fact  that  the  sulphuric  acid  itself  contains  both  of  these 
elements.  When  an  aqueous  solution  of  sodium  fluoride  NaF  is  electrolyzed,  the 
dissolved  substance  contains  neither  hydrogen  nor  oxygen.  Neither  the  sodium 
nor  the  fluorine  themselves  can  be  liberated,  however,  because  the  hydrogen  and 
oxygen  of  the  water  are  more  easily  discharged  and  set  free.  In  this  instance,  the 
salt  renders  the  liquid  a  good  conductor,  which  water  alone  is  not,  yet,  in  so  far 
as  decomposition  of  any  substance  is  concerned,  only  the  water  is  decomposed. 
This  is  a  straight  case  of  decomposing  water  by  electrical  energy. 

The  Other  Ways  of  Preparing  Hydrogen.  —  For  special  pur- 
poses, hydrogen  may  be  made  by  boiling  an  aqueous  solution  of 
sodium  hydroxide  with  aluminium  turnings,  when  sodium  aluminate 
is  formed:  Al  +  NaOH  +  H2O  ->  NaA102  +  3H;  also  by  heating 


122  INORGANIC  CHEMISTRY 

powdered  zinc  and  dry  sodium  hydroxide,  the  product  being  sodium 
zincate:  Zn  +  NaOH  +  H2O  -» NaHZn02  +  2H. 

Sources  of  the  Hydrogen  of  Commerce.  —  Zinc  is  too  expen- 
sive a  substance  to  use  in  the  preparation  of  hydrogen  in  large 
quantities  for  commercial  purposes.  We  realize  this  when  we  note 
that  33  parts  of  zinc  will  liberate  only  one  part  of  hydrogen,  so  that 
with  1  Ib.  of  zinc  we  obtain  only  one  half-ounce  of  the  gas.  Differ- 
ent sources  are  used  in  different  localities  and  countries. 

The  largest  supply  is  probably  obtained  as  a  byproduct  in  the 
electrolysis  of  an  aqueous  solution  of  common  salt  NaCl,  in  connec- 
tion with  the  manufacture  of  caustic  soda  (sodium  hydroxide,  q.v.). 
The  hydrogen  is  collected  and  compressed  in  steel  cylinders. 

In  some  circumstances,  the  method  of  passing  steam  over  heated 
iron  is  used  (p.  116). 

Another  plan  is  to  liquefy  water-gas  (q.v.),  a  mixture  of  hydrogen 
and  carbon  monoxide.  The  hydrogen  evaporates  much  the  more 
readily  of  the  two,  and  can  thus  be  separated.  This,  and  still  other 
processes  involve  substances  and  reactions  which  we  have  not  yet 
encountered  and  will  be  mentioned  at  the  appropriate  points. 

Preparation  of  Simple  Substances,  —  There  are  two  general 
sources  from  which  we  obtain  simple  substances.  If  the  element 
occurs  uncombined  in  nature,  as  sulphur  and  gold  do,  it  is  only 
necessary  to  free  it  from  foreign  materials  (impurities)  with  which 
it  is  mixed.  If  no  such  supply  exists,  or  if  the  purification  is  difficult, 
then  some  compound,  natural  or  artificial,  is  decomposed. 

The  liberation  of  an  element  from  combination  is  likewise  effected 
in  two  ways:  (1)  The  compound  may  be  forced  apart  by  the  appli- 
cation of  energy,  usually  in  the  form  of  heat  or  electrical  energy. 
Thus,  heat  was  employed  to  liberate  oxygen  (pp.  17,  82,  83).  Elec- 
trical energy,  however,  was  used  in  one  way  of  obtaining  hydrogen 
(p.  120),  since  readily  obtainable  compounds  which,  when  heated, 
give  hydrogen  as  the  only  gaseous  product,  are  unknown.  Elec- 
trical energy  may  be  used  also  to  liberate  any  metal.  (2)  The  other 
general  method  of  setting  an  element  free  is  to  offer  to  the  other  con- 
stituents of  the  compound  some  substance  with  which  they  will 
unite.  This  plan  is  used  in  liberating  hydrogen  from  acids  and 
from  water  (pp.  114-118). 

In  selecting  our  source,  we  are  naturally  influenced  by  the  cost 
of  the  material,  as  well  as  by  the  ease  of  the  process.  Thus,  gold 
oxide  yields  oxygen  by  the  application  of  very  little  heat,  but  it  is 


HYDROGEN 


123 


extremely  expensive.  Quicklime  is  very  cheap,  but  does  not  give 
up  its  oxygen  even  at  the  temperature  of  the  electric  arc.  As  a 
source  of  hydrogen,  any  acid  may  be  used,  theoretically,  but  the 
cheapest  and  commonest  acids  (hydrochloric  acid  and  sulphuric 
acid)  are  actually  employed. 

Purification  of  Gases.  —  Hydrogen  made  in  any  of  the  above 

ways  is  impure  (p.  6).  As  made  by  all  the  laboratory  methods 
mentioned,  a  good  deal  of  water 
vapor  is  mixed  with  it.  Other 
impurities,  like  hydrogen  sul- 
phide and  arsine,  come  from  the 
action  of  the  acid  on  foreign 
materials  in  the  zinc  (p.  119). 
Some  of  the  acid,  if  it  is  volatile, 
will  also  be  taken  over  with  the 
gas.  When  the  object  for  which 
the  gas  is  being  made  demands 
it,  we  must  know  what  the  im- 
purities to  be  expected  are,  and 
take  proper  means  of  removing 
them. 

Gases  are  freed  from  aque- 
ous vapor  by  means  of  calcium 
chloride  or  concentrated  sulphuric  acid,  which  greedily  absorb  mois- 
ture. The  former  is  used  in  granulated  form  in  straight 
or  bent  tubes  (Fig.  45).  The  latter  is  applied  by  sat- 
urating pieces  of  pumice-stone  with  the  acid  and  filling 
similar  tubes  with  the  fragments.  Or  the  acid  may  be 
placed  in  a  gas  washing  bottle  (Fig.  46).  For  extremely 
complete  drying,  a  tube  may  be  filled  with  phosphoric 
anhydride  sifted  upon  glass  beads  or  glass  wool.  Fore- 
thought must  be  used  to  avoid  a  drying  agent  which  will 
interact  with  gas.  The  longer  the  gas  remains  in  contact 
with  the  drying  agent,  the  more  perfect,  up  to  a  certain 
limit,  is  the  purification  effected.  In  all  .cases,  the  stream 
of  gas  must  pass  slowly. 
Particles  of  liquid  or  solid  matter  are  always  carried  along  by 
freshly  made  gases.  These  will  pass  with  the  gas  through  sulphuric 
acid  without  being  affected.  A  plug  of  cotton  or  of  glass  wool  in 
some  part  of  the  tubing  is  required  to  arrest  them. 


FIG.  45. 


FIG.  46. 


124  INORGANIC   CHEMISTRY 

Physical  Properties  of  Hydrogen.  —  Some  of  these  may  be 
given  in  tabular  form : 

Colorless  Grit,  temp.,  -234° 

Tasteless  Sp.  ht.  (gas,  const,  press.),  3.4 

Odorless  Boiling-point,  -252.5° 

Density  (air  =  1),  0.0695  Melting-point  (58  mm.),  -260° 

Wt.  of  1  L,  0.08987  g.  Sol'ty  in  Aq,  1.9  vols.  in  100  (14°) 

Air  is  14.5  times  heavier,  hence  the  gas  may  be  poured  upwards 
(Fig.  47)  and  is  used  for  filling  balloons.  A  liter  flask  filled  with  air 

requires  about  1.2  g.  to  be  added 
to  the  tare  to  restore  the  balance 
when  the  air  is  replaced  by  hy- 
drogen. Its  specific  heat  is  about 
seventeen  times  that  of  oxygen 
(0.2).  Its  thermal  conductivity 
is  greater  than  that  of  any  other 
elementary  gas.  Hence  a  wire, 
raised  to  incandescence  in  air  by 
means  of  an  electric  current,  can- 
not be  kept  at  a  red  heat,  even, 
by  the  same  current  in  hydrogen. 
This  is  due  to  the  fact  that  much 
of  the  heat  is  used,  not  merely 
FlG  47  to  increase  the  velocity  of  the 

molecules,  each  of  which  contains 

two  atoms  of  hydrogen,  but  also  to  decompose  them  into  atoms 
(see  Chap.  XII). 

Hydrogen  was  first  liquefied  in  visible  amounts  by  Dewar  (1898). 
The  liquid  is  colorless  and,  when  allowed  to  evaporate  rapidly  under 
reduced  pressure,  freezes  to  a  colorless  solid.  All  other  gases,  except 
helium,  solidify  easily  when  led  into  a  vessel  surrounded  by  liquid 
hydrogen. 

Hydrogen  is  absorbed,  for  the  most  part  in  a  purely  mechanical 
way,  by  many  metals.  Heated  iron  will  take  up  19  times  its  volume 
of  hydrogen.  Under  similar  conditions  gold  takes  up  46  volumes, 
platinum  in  fine  powder  50  volumes,  palladium  502  volumes,  and 
silver  none.  The  maximum  absorbed  by  palladium  under  favorable 
cond;tions  is  873  volumes.  It  is  still  a  question  whether,  in  the 
case  of  palladium,  a  part  of  the  gas  is  not  in  combination. 


HYDROGEN 


125 


Diffusion.  —  If  a  volume  of  gas  is  enclosed  at  one  end  of  a  cyl- 
inder, the  rest  of  which  is  entirely  empty,  and  is  suddenly  released 
from  this  confinement,  it  spreads  with  extreme  speed  so  as  to  occupy 
the  whole  of  the  cylinder  to  an  equal  degree.  This  spreading  is  not 
an  effect  of  gravitation,  since  it  takes  place  upwards  or  downwards 
with  equal  celerity.  The  same  phenomenon  is  observed  when,  in 
everyday  life,  a  bottle  of  scent  is  opened.  The  vapor,  on  escaping, 
begins  to  penetrate  in  all  directions  through  the  room,  showing  its 
presence  by  its  odor.  The  motion,  as  this  instance  shows,  takes 
place  through  a  space  occupied  by  another  gas  more  slowly  than, 
but  just  as  surely  as,  when  the  space  is  empty.  The  material  of 
gases  has  in  fact  an  independent  power  of  locomotion.  The  result- 
ing phenomenon  we  call  diffusion.  It  is  constant  in  rate  for  each 
gas  under  like  conditions,  and  hydrogen  has  the  greatest  speed  of 
diffusion  of  all  the  gases. 

The  interdiffusion  of  gases  and  the  absence  of  gravity  effect  may  be  shown 
simultaneously.  A  jar  of  air  is  inverted  and  placed  mouth  to  mouth 
with  a  jar  filled  with  carbon  dioxide  (Fig.  48).  After  a  few  minutes, 
and  in  spite  of  the  fact  that  carbon  dioxide,  measured  in  bulk,  is 
one-half  heavier  than  air,  limewater  will  show  the  presence  of  car- 
bon dioxide  in  the  upper  jar.  The  phenomena  of  diffusion  must 
not  be  confused  with  cases  like  the  pouring  of  hydrogen 
upward  to  displace  air  in  an  inverted  jar.  In  this  case 
the  gas  flows  en  masse,  and  the  gravity  effect  is  the 
very  one  on  which  we  depend  for  the  success  of  the 
experiment.  It  is  when  hydrogen  scatters  itself  in  a 
somewhat  slower  way,  and  downward  and  sideways  as  well  as 
upward,  that  we  have  diffusion.  The  word  indicates  the  scattering 
rather  than  the  flowing  nature  of  the  phenomenon. 

The  different  rates  of  diffusion  of  different  gases  are 
easily  shown  by  comparing  their  several  speeds  with  that 
of  air,  when  both  pass  through  a  wall  of  unglazed,  porous 
porcelain. 

The  cylinder  of  porous  porcelain  A  (Fig.  49)  contains 
air  and  is  connected  with  a  wide  tube  which  dips  beneath 
the  surface  of  the  water.  When  a  cylinder  H  containing 
hydrogen  is  brought  over  it,  rapid  escape  of  gas  takes 
place  through  the  water,  showing  that  a  rise  in  pressure 
has  taken  place  inside  the  porous  vessel.  Before  the 
cylinder  of  hydrogen  approached  it,  the  air  was  moving  both  out- 
wards and  inwards  through  the  porcelain,  but,  being  the  same 


FIG.  48. 


FIG.  49. 


126  INORGANIC  CHEMISTRY 

air,  the  speed  of  motion  was  equal  in  both  directions,  and  therefore 
the  pressure  inside  was  not  affected.  It  is  important  to  note  that 
there  was  at  no  time  rest,  there  was  simply  equal  motion  in  both 
directions.  When  the  hydrogen  atmosphere  surrounded  the  cylin- 
der, the  hydrogen  gas  moved  more  rapidly  into  the  cylinder  than 
the  air  inside  could  move  out,  and  hence  an  excess  of  pressure  quickly 
arose  in  the  interior. 

Exact  measurement  shows  that  the  lighter  a  gas  is  in  bulk,  the 
faster  its  parts  move  by  diffusion  in  any  direction.  The  rate  is 
inversely  proportional  to  the  square  root  of  the  density  of  the  gas. 
Thus,  for  hydrogen  and  air  it  is  in  the  ratio  Vl.293  :  Vo.08987,  or 
3.8  :  1. 

Chemical    Properties    of   Hydrogen.  —  Hydrogen,  delivered 
from  a  jet,  burns  in  air  or  pure  oxygen.     A  cold  vessel,  held  over  the 
almost  invisible  blue  flame  (Fig.  50),  condenses  to  droplets  of  water 
the  steam  that  is  produced.     When  hydrogen 
and  oxygen  are  mingled  in  a  suitable  burner 
(Fig.  51),  and  the  flame  is  allowed  to  play  on 
a  piece  of  quicklime,  the  latter  becomes  white- 
hot  at  the  spot  where   the  flame  meets  it. 
This  result  is  called  a  calcium  light  or  lime 
_  light.     Platinum  melts  in  this  flame  easily.     In 

Ja  closed  space  it  produces  a*  temperature  of 
over  2500°. 

FIG  50  When  oxygen  and  hydrogen  are  mixed  in 

a  glass  vessel,  the  chemical  action  is  very  slow 
at  ordinary  temperatures,  no  perceptible  amount  of  union 
occurring  in  a  period  of  five  years.  If  the  mixture  is  sealed  up  and 
kept  at  300°,  after  several  days  a  small  part  is  found  to  have  com- 
bined to  form  water.  At  518°,  hours  are  required  before  the  union 
is  complete.  At  600°  the  interaction  is  rapid,  but  not  explosive. 
At  700°  the  combination  is  almost  instantaneous.  Hence  contact 
with  a  body  at  a  bright-red  heat  (p.  94)  is  required  actually  to 
explode  the  mixture. 

These  facts  illustrate  the  effect  of  temperature  on  the  speed  of 
chemical  changes  (p.  93).  A  rough  calculation  shows  that,  since 
interactions  lower  their  speed  to  half  its  value  for  every  depression  of 
10°  in  temperature,  at  ordinary  temperatures  this  union  can  hardly 
make  easily  perceptible  progress  in  less  than  a  thousand  million 
years.  This  effect  of  temperature,  therefore,  accounts  for  the  ap- 
parent absence  of  action  in  the  cold  gases. 


HYDROGEN  127 

Finely  divided  platinum,*  when  held  in  the  cold  mixture,  hastens 
the  otherwise  vanishingly  slow  union  in  the  part  of  the  gases  in  con- 
tact with  it.  The  heat  of  their  union  raises  the  temperature  of  the 
platinum  to  a  white  heat  and  this  causes  explosion  of  the  whole  mass 
of  gas.  The  platinum  is  simply  a  contact  agent  (p.  97)  and  remains 
itself  unaffected.  The  same  explanation  applies  to  self-lighting  gas 
jets  and  to  pocket  cigar  lighters. 

Hydrogen  unites  directly  with  a  minority  only  of  the  simple  sub- 
stances. It  combines  rapidly  with  oxygen,  chlorine,  fluorine,  and 
lithium,  and  more  slowly  with  a  few  others. 

Hydrogen  also  unites  with  oxygen  and  chlorine,  even  when  these 
elements  are  already  combined  with  certain  of  the  metals.  Thus, 
when  one  of  the  oxides  of  copper  or  of  iron  is  heated  in  a  tube  through 
which  hydrogen  flows,  the  latter  combines  with  the  oxygen  to  form 
water,  and  the  metal  is  liberated : 

Skeleton:  CuO   +H->H2O  +  Cu. 

Skeleton:  Fe3O4  +  H  ->  H2O  +  Fe. 

We  observe  that,   for  each  atomic  weight  of  oxygen,  2H  will  be 
required,  and  amend  the  equations  thus: 

Balanced:  CuO  +  2H-»H20  +  Cu, 

Skeleton:  Fe3O4  +  8H  ->  4H2O  +  Fe. 

Then  we  take  the  amount  of  iron  produced  equal  to  that  taken : 
Balanced:  Fe304  +  8H  -*  4H2O  +  3Fe. ' 

These  interactions  are  classed,  by  mechanism,  as  displacements 
(p.  21).  In  describing  them  the  chemist  would  also  say  that  the 
hydrogen  has  been  oxidized  and  that  the  oxide  of  the  metal  has  been 
reduced  (p.  91). 

An  Inapt  Use  of  the  Word  "  Affinity  "  in  Explanation  of 
Chemical  Actions.  —  Speaking  without  reflection,  one  might  be 
tempted  to  "explain"  actions  like  the  reduction  of  magnetic  oxide 
of  iron,  just  mentioned,  by  saying  that  the  hydrogen  has  a  greater 
tendency  to  unite  with  oxygen,  or  has  a  greater  affinity  for  it,  than 
has  iron,  and  therefore  removes  the  oxygen  from  combination  with 
the  latter.  Plausible  as  this  statement  seems,  it  would  be  in  many 

*  The  most  convenient  form  is  obtained  by  dipping  asbestos  in  a  solution  of 
chloroplatinic  acid,  and  heating  it  in  the  blast-lamp.  The  fibers  are  covered  with 
a  thin  film  of  the  metal:  HaPtCl«  — >  Pt  +  4C1  +  2HC1. 


128  INORGANIC  CHEMISTRY 

cases,  as  here,  quite  incorrect.  Under  the  modes  of  preparing  hydro- 
gen, we  spoke  of  the  action  of  steam  upon  iron  (p.  116),  and  gave 
the  equation:  3Fe  +  4H2O  -»  Fe3O4  +  8H.  To  use  consistently 
this  handy  method  of  explaining  chemical  change  by  the  help  of  the 
word  " affinity,"  we  should  have  to  say  here  that  the  hydrogen  has 
a  less  affinity  for  the  oxygen  than  has  iron,  and  therefore  hydrogen 
is  set  free  and  oxide  of  iron  is  formed.  It  will  be  seen  that  this 
statement  is  in  direct  contradiction  to  the  one  made  above.  Both 
cannot  be  true.  The  fact  is  that  both  are  based  upon  an  assumption 
which  is  incorrect  —  the  assumption,  namely,  that  the  displacement 
of  one  element  by  another  is  always  an  evidence  of  the  greater 
affinity  of  the  latter.  The  correct  explanation  of  seemingly  contra- 
dictory actions  like  these  will  be  given  later  (see  Hydrogen  chloride). 

The  action  of  catalytic  agents  is  itself  a  refutation  of  this  blundering  assump- 
tion. Putting  a  little  platinum  in  a  mixture  of  oxygen  and  hydrogen  cannot  add 
to  the  energy  contained  in  these  substances,  and  cannot  therefore  increase  their 
intrinsic  tendencies  to  unite.  Yet  in  its  presence  an  almost  nonexistent  action 
becomes  suddenly  explosively  violent.  There  are  other  factors,  often  far  more 
potent  than  affinity,  which  determine  the  direction  and  speed  of  many  chemical 
changes  (see  Chemical  equilibrium). 

In  this  connection,  it  is  worth  noting  that,  while  increasing  the  speed  of  a  train 
or  a  ship  requires  a  great  addition  to  the  energy  expended,  and  is  verj  costly,  in- 
creasing the  speed  of  a  chemical  change  requires  the  expenditure  of  no  energy 
whatever.  The  employment  of  a  couple  (p.  119)  or  a  catalytic  agent  adds  nothing 
to  the  energy  the  separate  bodies  possessed  before  they  were  mixed.  And  the  cata- 
lytic agent  is  recovered  unchanged  and  as  efficient  as  ever  at  the  end.  Theoreti- 
cally, therefore,  these  agencies  cost  nothing.  The  increased  speed  in  the  forma- 
tion of  the  products  is  obtained  gratis.  The  contact  method  of  making  sulphuric 
acid  (q.v.)  illustrates  the  way  in  which  commerce  has  taken  advantage  of  this  fact. 

The  Speed  of  Chemical  Actions:  a  Means  of  Measuring 
Activity.  —  The  speed  of  a  chemical  action  is  measured  by  the 
number  of  atomic  or  formula  weights  of  the  substance  undergoing 
change  in  a  given  time.  Now,  one  means  of  measuring  the  relative 
chemical  activities  of  several  substances,  is  to  observe  the  speed 
with  which  they  undergo  the  same  chemical  change  (p.  37).  Thus 
we  may  compare  the  activities  of  the  various  metals  by  allowing 
them  separately  to  interact  with  hydrochloric  acid  and  collecting 
and  measuring  the  hydrogen  liberated  per  minute  by  each.  It  will 
be  seen,  even  in  the  roughest  experiment,  that  magnesium  is  thus 
much  more  active  than  zinc.  The  comparison  must  be  made  with 
such  precautionsj  however,  as  will  make  it  certain  that  the  condi- 


HYDROGEN 


129 


tions  under  which  the  several  metals  act  are  all  alike.  Thus,  in 
spite  of  the  heat  evolved  by  the  action,  means  must  be  used,  by 
suitable  cooling,  to  keep  the  temperature  at  some  fixed  point  during 
all  the  experiments,  for  all  actions  become  more  rapid  when  the 
temperature  rises  (p.  93).  Again,  pure  materials  must  be  used,  as 
an  impurity  in  one  metal  might  act  as  a  contact  agent  and  modify 
the  natural  speed  of  the  action  (p.  119).  Still  again,  the  pieces  of 
the  various  metals  must  be  arranged  so  that  equal  surfaces  are  ex- 
posed to  the  acid  in  each  case.  Equal  weights  of  zinc  will  finally 
give  equal  weights  of  hydrogen;  but  if  one  of  them  is  in  the  form  of 
foil  while  the  other  is  a  cylinder,  the  former,  although  it  will  not  last 
so  long,  will  give  more  hydrogen  per  minute.  Finally,  the  portions 
of  hydrochloric  acid  must  contain  the  same  percentage  of  hydrogen 
chloride  in  each  case,  for  the  metal  will  secure  the  acid  it  needs  with 
less  delay  in  a  more  concentrated  solution  than 
in  a  less  concentrated  solution,  and  in  the  former 
case  will  therefore  displace  hydrogen  more  rapidly. 
When  these  and  other  precautions  have  been 
taken,  a  true  comparison  of  the  relative  activities 
of  the  metals  with  respect  to  this  particular  ac- 
tion may  be  made.  It  is  found  that  the  order  in 
which  this  comparison  places  the  metals  is  much 
the  same  as  that  in  which  they  are  placed  by  a 
study  of  other  similar  actions.  This  is  natural, 
since  we  are  really  comparing,  in  each  case,  the 
amount  of  free,  internal  energy  in  each  metal. 
A  single  table  suffices,  therefore,  for  all  purposes 
(see  next  section). 


ORDER  OP 
ACTIVITY. 
METALS. 


The  Order  of  Activity  of  the  Metals. — We 

employ  metals  so  frequently  in  chemistry,  that  we 

must  at  once  become  familiar  with  the  key  to  the 

main  differences  in  their  behavior.     The  order  of 

their  activity  explains  these  differences,  as  well 

as  many  other  facts.     In  the  adjoining  list,  the 

most  active  metals  are  at  the  top.     Hydrogen  is 

not  a  metal,  but  is  included  because  chemically 

it  resembles  the  metals.     All  the  metals  above  hydrogen  displace 

this  element  from  dilute  acids  (and  from  water),  while  those  below 

it  do  not. 

The  first  displaces  the  hydrogen  from  water  violently,  the  second 


Potassium 

Sodium 

Calcium 

Magnesium 

Aluminium 

Manganese 

Zinc 

Chromium 

Iron 

Nickel 

Lead 

Tin 

Hydrogen 

Copper 

Bismuth 

Antimony 

Mercury 

Silver 

Platinum 

Gold 


130  INORGANIC  CHEMISTRY 

less  vigorously.  Magnesium  barely  acts  on  boiling  water,  but,  like 
iron,  acts  on  superheated  steam.  Zinc  liberates  hydrogen  with 
reasonable  vigor  from  dilute  acids,  lead  rather  feebly,  and  copper 
and  those  following  not  at  all. 

Other  facts  are  explained  by  the  table.  Thus,  when  the  metals 
are  heated  in  pure  oxygen,  the  last  three  do  not  combine.  Those 
above  silver  do  unite  with  oxygen  —  mercury  rather  slowly  and  the 
others  more  and  more  energetically  as  we  ascend  the  list.  Again,  if 
we  take  the  oxides  of  the  metals,  we  find  that  those  of  the  metals 
up  to  and  including  mercury  lose  all  their  oxygen  when  simply 
heated.  If  we  heat  the  oxides,  and  also  lead  hydrogen  over  them, 
the  oxygen  is  easily  removed  from  all  the  oxides  up  to  and  including 
those  of  iron,  leaving  in  each  case  the  metal.  Thus,  in  general,  the 
more  active  metals  form  the  most  stable  compounds. 

The  metals  following  hydrogen  are  the  ones  which  are  found  in 
nature  in  large  amounts  in  the  free  condition.  Those  preceding 
hydrogen,  if  liberated,  since  they  all  displace  hydrogen  from  acids, 
would  act  slowly  on  natural  waters  containing  carbonic  acid,  and  so 
would  pass  into  combination.  Hence  those  preceding  hydrogen 
are  found  free  only  in  exceptional  cases,  such,  for  example,  as  the 
metallic  iron  of  which  meteorites  often  consist. 

VALENCE 

Equivalence  and  Valence.  —  If  the  equations  showing  dis- 
placement of  hydrogen  by  a  metal  be  now  reexamined,  a  peculiarity 
will  be  noticed  which  we  have  thus  far  omitted  to  note.  When 
sodium  (p.  115)  and  calcium  (p.  116)  act  upon  water,  one  atomic 
weight  (or  atom)  of  sodium  displaces  one  atomic  weight  of  hydrogen, 
but  one  atomic  weight  of  calcium  displaces  twice  as  much  hydrogen. 
Again,  one  atom  of  zinc  (p.  118)  displaces  two  atoms  of  hydrogen, 
but  one  atom  of  aluminium  displaces  three.  Assuming,  for  simplic- 
ity, that  we  allow  three  of  these  metals  all  to  act  upon  dilute  hydro- 
chloric acid,  the  equations  are: 

Na  +  HCl  ->NaCl  +  H 
Ca  +  2HCl->CaCl2  +  2H 
A1  +  3HC1-*A1C13  +  3H. 

Interpreting  this,  we  perceive  that  the  atom  of  aluminium,  for 
example,  displaces  3H,  because  it  is  able  to  combine  with  3CI,  and  so 
incidentally  liberates  the  hydrogen  formerly  united  with  3C1.  The 


HYDROGEN  131 

atom  of  sodium,  however,  can  unite  with  only  1C1,  and  so  releases 
only  1H.  Now  this  is  not  a  rule  confined  to  these  reactions,  but 
represents  a  general  chemical  property  of  the  atomic  weight  of  each 
element,  and  a  property  which  we  shall  find  most  useful. 

The  atom  of  aluminium  releases  3H  because  it  can  take  the 
place  of  three  atoms  of  hydrogen  in  chemical  combination  (and 
hold  3C1).  The  atomic  weight  of  aluminium  is  said  to  be  equivalent 
to  (equal  in  chemical  value  to)  three  atomic  weights  of  hydrogen. 
Since  it  combines  with  3  atomic  weights  of  chlorine,  it  is  also  con- 
sidered to  be  equivalent  to  three  atomic  weights  of  this  element. 

The  chemical  property  referred  to  is  called  valence.  The  valence 
of  an  atomic  weight  of  hydrogen  or  of  chlorine  is  the  unit.  An 
atomic  weight  of  sodium  is  said  to  be  univalent,  one  of  calcium 
bivalent,  one  of  aluminium  trivalent.  The  formula  H2O  shows  the 
atomic  weight  of  oxygen  to  be  bivalent,  because  it  unites  with  two 
atomic  weights  of  hydrogen.  Apparently,  the  atomic  weight  (or 
atom)  of  each  element  has  a  fixed  capacity  for  combining  with  not 
more  than  a  certain  number  of  atomic  weights  (or  atoms)  of  some 
other  element. 

Marking  the  Valence.  —  Until  we  have  become  familiar  with 
the  valence  of  each  element,  it  is  advisable  to  mark  the  valences  in  a 
special  way:  Na1,  CarrrAlm,  Ou,  Zn11,  Cl1. 

As  we  should  expect,  a  bivalent  atom  can  combine  with  two 
univalent  atoms,  or  with  one  bivalent  atom,  and  so  forth.  Thus  we 
have  the  compounds  of  oxygen:  Na/O11,  Ca1^11,  Al2m03n,  Zn^O11, 

cvo11. 

The  rule  is  that  the  quantities  of  two  elements  which  combine 
must  have  equal  total  combining  capacities  —  i.e.,  identical  total 
valence.  Thus,  Ca11  has  the  valence  two,  and  so  does  OK.  Again, 
Al2m  has  a  total  valence  of  2  X  3  (=  6)  and  so  has  03"  (3X2  =  6). 

Frequently  the  valence  is  marked  by  means  of  lines,  the  number 
of  lines  pointing  towards  a  symbol  indicating  the  valence  of  the  atom 
it  represents: 


Na-Cl      Ca  Ca  =O      A1-C1      0  =  Al  -  O  -  Al  =  O 

XC1  XC1 

Definition.  —  The  valence  of  an  element  is  a  number  represent- 
ing the  capacity  of  its  atomic  weight  to  combine  with,  or  displace, 


132  INORGANIC  CHEMISTRY 

atomic  weights  of  other  elements,  the  unit  of  such  capacity  being 
that  of  one  atomic  weight  of  hydrogen  or  chlorine. 

Stated  otherwise,  the  valence  of  the  atomic  weight  (or  atom)  of 
an  element  is  the  number  of  atomic  weights  (or  atoms)  of  hydrogen, 
or  of  some  other  univalent  element,  which  the  atomic  weight  (or 
atom)  of  the  given  element  combines  with  or  displaces. 

Valence  of  Radicals.  —  What  we  have  said  applies  to  com- 
pounds of  not  more  than  two  elements  —  so-called  binary  com- 
pounds. We  cannot  with  certainty  tell  the  valences  in  a  compound 
of  three  or  more  elements,  like  H2SO4.  But  we  have  seen  that  the 
acids  behave  as  if  composed  of  two  radicals  (p.  117) :  H(C1),  H2(SO4), 
that  is,  of  two  groups  which  move  as  wholes  in  chemical  reactions. 
Thus,  in  the  interaction  of  zinc  with  dilute  sulphuric  acid: 

Zn  +  H2SO4  -^  ZnS04  +  2H 

the  group  SO4  passes  as  a  whole  from  combination  with  2H  into 
combination  with  Zn.  Hence  we  can  assign  a  valence  to  a  com- 
pound radical  as  a  whole.  Thus,  (SO4)n  is  evidently  bivalent,  as  a 
whole,  because  it  is  united  with  2H1.  Na(OH)  and  Ca(OH)2  show 
the  radical  hydroxyl  (OH)1  to  be  univalent. 

It  is  to  preserve  the  identity  of  the  radicals,  and  to  make  them 
easily  recognizable,  that  we  write  them  in  ^brackets  and  place  the 
coefficient  outside,  as  Ca(OH)2  and  A12(S04)3,  instead  of  using  the 
forms  Ca02H2,  Al2S3Oi2,  and  so  forth.  The  bracketed  forms  show 
more  clearly  which  radicals  are  present.  Since  substances  like  these 
commonly  interact  by  double  decomposition  (p.  20)  or  displacement 
(p.  18),  as  if  the  radicals  were  single  elements,  this  mode  of  writing 
the  formulae  enables  us  to  use  the  valence  values,  and  so  readily  to 
write  equations  for  such  actions.  This  artifice,  justified  as  it  is  by 
the  mode  of  interaction,  reduces  many  substances  containing  three 
or  more  elements  to  binary  compounds.  In  writing  formulae  of 
inorganic  compounds,  we  usually  place  the  positive  radical  (p.  117) 
in  front  and  the  negative  radical  after  it. 

Use  in  Making  Formulse  and  Equations.  —  The  chief  use  of 
the  conception  of  valence  is  the  very  practical  one  of  enabling  us  to 
write  formulae.  In  making  equations  we  constantly  need  to  know 
whether  the  chloride  of  an  element,  say  magnesium,  is  MgCl,  or 
MgCl2,  or  MgCl3,  or  MgCl4,  etc.,  and  whether  its  sulphate  is  MgSO4, 
Or  Mg2S(>4,  or  some  other  combination  of  the  symbols.  To  answer 


HYDROGEN  133 

questions  like  this  it  is  not  necessary  to  know  the  formula  of  every 
compound  of  each  element:  the  apparent  disorder  of  these  numbers 
can  be  reduced  to  rule,  and  the  reader  should  endeavor  thoroughly 
to  master  the  rule  before  going  farther. 

Thus,  suppose  that  we  require  the  formula  of  aluminium  hydrox- 
ide. Up  to  this  point,  we  should  have  been  compelled  to  look  for 
it  in  a  book.  And  if,  later,  we  needed  the  formula  of  aluminium 
sulphate,  we  should  have  had  to  look  that  up,  separately,  also. 
But  now,  all  we  need  is  to  know  the  valence  of  aluminium  Alm,  of 
the  hydroxyl  radical  (OH)1,  and  of  the  sulphate  radical  (SOJ11. 
Making  the  total  valences  in  each  half  of  each  compound  alike,  we 
write  the  formulse  Alm  (OH)3I,  A12III(SO4)311. 

The  reader  must  make  a  special  effort  to  note  the  valences  of 
each  element  and  radical,  and  always  to  use  them  in  making  formulae. 
If  a  formula  is  written  from  memory,  the  valences  must  be  checked, 
to  make  sure  that  the  formula  is  correct. 

How  to  Learn  the  Valence  of  an  Element.  —  To  find  out  the 
valence  of  an  element,  we  must  obtain  the  formula  of  one  simple 
compound  of  the  element,  containing  another  element  of  known 
valence.  Thus,  what  is  the  valence  of  carbon?  Its  oxide  is  CC>2. 
The  total  valence  of  oxygen  here  is  2X2  =  4.  Carbon  C1V  is 
therefore  quadrivalent.  Hence  its  chloride  must  be  C^CU1  (carbon 
tetrachloride) ,  and  its  compound  with  hydrogen  C^H/  (methane, 
composing  a  large  part  of  natural  gas).  When  carbon  combines 
with  a  trivalent  element,  equi-valent  amounts  of  each  element  must 
be  used,  as  in  A14UI  C31V  (aluminium  carbide),  where  Alf1  and  CJ* 
contain  3  X  4,  or  12  units  of  valence  each. 

Again,  when  we  know  the  formula  of  sodium  iodide  to  be  Na1!, 
or  that  of  hydrogen  iodide  to  be  H1!,  we  infer  that  iodine  is  univalent. 
The  formula  of  silica  (sand)  Si02IX  shows  silicon  to  be  quadrivalent, 
and  indicates  that  the  chloride  must  be  SiCl4.  Similarly  the  formula 
of  calcium  carbonate  CanCO3  shows  that  the  radical  C03,  which  is 
common  to  all  carbonates,  must  be  bivalent. 

The  chemist  does  not  memorize  the  valences  themselves;  he  re- 
covers the  valence  of  an  element  or  radical,  when  needed,  by  recalling 
the  formula  of  a  substance  containing  this  element  or  radical  in  com- 
bination with  a  more  familiar  element  or  radical,  such  as  Cl1  or  H1. 

Elements  with  More  than  One  Valence.  —  The  rule  of  va- 
lence is  somewhat  complicated  by  the  fact  that  many  elements  show 


134  INORGANIC  CHEMISTRY 

more  than  one  valence.  In  other  words,  the  combining  capacity 
of  an  atomic  weight  of  such  an  element  may  have  two  (or  even 
more)  valences,  according  to  the  conditions  under  which  the  action 
takes  place.  This  is  as  much  as  to  say  that  an  atomic  weight  of 
such  an  element  may  form  stable  compounds  with  two,  or  even 
more  different  numbers  of  equivalents  of  another  element.  This 
fact  has  already  been  mentioned,  for  it  is  implied  in  the  law  of  mul- 
tiple proportions  (p.  58). 

Thus,  we  have  encountered  two  chlorides  of  iron,  ferrous  chlo- 
ride FenCl2I  and  ferric  chloride  Fe111^1.  We  have,  in  fact,  two 
complete  series  of  compounds  of  iron,  such  as: 

Bivalent  (Ferrous):        FeCl2,  FeO,  FeSO4. 
Trivalent  (Ferric):         FeCl3,  Fe2O3,  Fe2(SO4)3. 

When  an  element  forms  two  such  series  of  compounds,  we  always 
call  particular  attention  to  the  fact. 

As  a  rule,  an  element  passes  from  one  form  of  combination  to 
another  without  change  of  valence.  But  compounds  of  elements 
like  tin  can  also  undergo  changes  in  course  of  which  the  valence 
alters.  Cases  of  this  kind  will  be  considered  when  they  arise  (see 
Preparation  of  chlorine). 

Exceptional  Valences.  —  Some  elements  show  an  exceptional 
valence  in  one  compound.  The  valences  shown  in  series  of  com- 
pounds are  the  important  ones,  and  the  exceptions  need  not  par- 
ticularly concern  us.  Thus,  in  addition  to  the  oxides  FeO  and  Fe203, 
iron  gives  the  magnetic  oxide  Fe3O4,  where  the  valence  of  iron  ap- 
pears not  to  be  a  whole  number,  but  f  or  2§.  The  valence  is  made 
regular  by  supposing  the  oxide  to  be  a  compound  of  the  other  two 
oxides,  as  if  the  formula  were  FeO,Fe203. 

Nomenclature.  —  The  names  of  compounds  containing  only 
two  elements  (the  true  binary  compounds)  end  in  ide.  Such  are 
the  oxides,  as  ferric  oxide  Fe2O3;  the  carbides,  as  aluminium  carbide 
ALiC3;  the  chlorides,  as  sodium  chloride  NaCl;  the  sulphides,  as 
ferrous  sulphide  FeS,  etc.  This  applies  also  to  compounds  con- 
taining positive  compound  radicals,  as  ammonium  chloride  (NH4)C1. 

When  an  element  forms  two  (or  more)  compounds  with  another 
element,  they  are  frequently  distinguished  thus:  carbon  dioxide 
CO2,  carbon  monoxide  CO;  phosphorus  pentoxide  P^O^,  phosphorus 
Zrioxide  P20g 


HYDROGEN  135 

To  distinguish  two  compounds  of  the  same  elements,  another 
plan  is  also  used;  ferrous  chloride  FeCU,  feme  chloride  FeCla;  mer- 
curous  oxide  Hg20,  mercuric  oxide  HgO.  The  suffix  ous  indicates 
that  the  metal  is  combined  with  the  smaller  proportion  of  the 
negative  element,  and  ic  that  it  is  combined  with  the  larger  pro- 
portion. 

The  tendency  —  although  not  a  universal  rule  —  is  to  use  the 
latter  plan  with  compounds  containing  a  metal  and  the  former  with 
compounds  containing  only  non-metals. 

Valence  and  Equivalent  Weights:   A  Different  Viewpoint. 

—  In  the  foregoing  discussion  of  valence,  we  have  more  than  once 
used  the  word  "equivalent."  If  the  method  by  which  the  atomic 
weights  were  derived  from  equivalents  (p.  63)  be  now  reexamined, 
a  different,  and  instructive  view  of  the  nature  of  valence  will  be 
obtained.  It  was  found,  for  example,  that  9.03  parts  of  aluminium 
(p.  63)  combined  with  the  equivalent  weights  of  the  other  elements, 
and  therefore  with  35.46  parts  of  chlorine.  If  this  weight  of  alu- 
minium had  been  accepted  as  the  final  unit  (the  atomic  weight), 
then  it  would  have  been  represented  by  the  symbol  Al  and,  since  Cl 
stands  for  35.46  parts  of  chlorine,  the  formula  of  the  chloride  would 
have  been  A1C1.  In  point  of  fact,  however,  a  number  three  times 
as  large  as  the  equivalent,  namely  27.1,  was  chosen  as  the  atomic 
weight  of  aluminium,  and  the  symbol  Al  stands  for  this  triple  quan- 
tity. If  the  equivalent  of  chlorine  had  also  been  tripled  in  making 
its  atomic  weight,  the  amounts  represented  by  the  symbols  would 
still  have  been  chemically  equivalent  and  the  formula  would  still 
have  been  A1C1.  But  the  equivalent  of  chlorine  was  left  unaltered. 
Hence,  to  get  the  equivalent  amounts  (i.e.,  the  actual  combining 
quantities)  of  the  two  elements,  we  must  have  3C1  with  1A1.  The 
formula  is  thus  A1C13.  Now,  it  is 'evident  that  this  tripling  of  the 
equivalent  of  aluminium  will  affect  the  formulae  of  all  its  compounds. 
Whenever  it  is  combined  with  an  element  which,  like  chlorine,  has 
identical  equivalent  and  atomic  weights,  the  formula  of  the  com- 
pound will  be  of  the  form  A1X3.  In  accordance  with  this  we  have 
the  bromide  AlBr3.  In  making  the  formulae  of  compounds  of  alu- 
minium, the  chief  thing  to  be  kept  in  mind,  therefore,  is  the  fact 
that  its  atomic  weight  contains  three  equivalents  and  always  com- 
bines with  three  equivalents  of  another  element.  This  fact  we  state 
by  saying  that  the  valence  of  the  atomic  weight  of  aluminium  is 
three,  or  simply  that  the  element  aluminium  is  trivalent. 


136  INORGANIC  CHEMISTRY 

Similarly,  the  equivalent  of  tin  is  59.5  and  its  atomic  weight  is 
119.  This  atomic  weight  therefore  contains  two  equivalents  of  tin 
and  combines  with  two  equivalents  of  any  other  element.  Hence, 
the  formula  of  a  compound  of  tin  with  an  element  of  the  chlorine 
class  will  be  SnX2.  Thus  tin  is  bivalent.  In  like  manner  the  equiva- 
lent of  sodium  is  23,  and  this  number  was  not  altered  in  making  the 
atomic  weight.  Hence,  the  symbol  Na  stands  for  one  equivalent, 
and  the  formula  of  the  compound  with  chlorine  is  NaCl.  Elements 
whose  atomic  weights  are  identical  with  their  equivalents  are  de- 
scribed as  univalent. 

Thus  the  valence  of  an  element  may  be  defined  as  the  number  of 
equivalent  weights  contained  in  its  atomic  weight.  Arithmetically 
it  is  the  integer  by  which  the  equivalent  weight  was  multiplied  in 
forming  the  atomic  weight. 

The  above  mode  of  handling  valence  is  based  upon  the  notion  of1 
combination  in  equivalent  proportions.  Another  variety  of  chemical 
change,  namely  displacement  (p.  18),  is  often  of  assistance  in  ena- 
bling us  to  determine  the  valence  of  an  element.  It  will  be  noted 
that  when  Al  acted  upon  hydrochloric  acid  (p.  118)  and  combined 
with  3C1,  it  necessarily  displaced  the  3H  with  which  the  3C1  was 
formerly  united.  It  was  equivalent  to  3H  for  the  purpose  of  holding 
3C1  in  combination.  It  is  from  this  aspect  of  the  relation  that  the 
word  "valence"  comes.  Al  is  equi-valent  to  3H,  and,  H  having  the 
unit  valence,  Al  is  trivalent.  Similarly,  since  one  atomic  weight  of 
zinc,  represented  by  the  symbol  Zn,  displaces  2H  (p.  118)  zinc  must 
be  bivalent.  Combining  this  with  the  former  conception,  we  reach 
a  definition  of  the  valence  of  an  element:  The  valence  of  an  element  is 
the  number  of  equivalent  weights  which  its  atomic  weight  contains  and 
is  therefore  the  number  of  equivalent  weights  of  another  element  or 
radical  which  its  atomic  weight  is  able  to  combine  with  or  displace. 

It  will  be  seen  that  the  equivalent  weight  can  always  be  found 
by  a  quantitative  experiment.  It  is  also  evident  that  it  is  equal  to 
the  atomic  weight  divided  by  the  valence.  It  is  likewise  clear  that 
the  equivalent  weight  of  an  element,  multiplied  by  the  valence  of 
that  element,  is  equal  to  the  atomic  weight.  The  conception  of 
equivalent  weight  finds  application  in  several  connections  in  chem- 
istry (see  Normal  Solutions  and  Faraday's  Law).. 

As  we  have  seen  (p.  134),  the  regular  valence  of  an  element  cannot  be  learned 
by  examining  the  composition  of  a  compound  chosen  at  random.  Thus  FeS, 
H2S,  HgS,  and  other  compounds  show  sulphur  to  be  bivalent.  There  is  also  a 
series  in  which  sulphur  is  sexivalent,  as  in  SOs.  But  the  compound  S2O3,  in  which 


HYDROGEN  137 

sulphur  appears  to  be  trivalent,  is  an  isolated  case.  Again,  FeO,  FeS,  FeCl2 
show  iron  to  be  bivalent,  and  FeCls,  Fe2(SO4)3,  etc.,  show  it  to  be  also  trivalent. 
But  Fe3O4,  the  magnetic  oxide,  is  an  exception.  Valence  has  to  do  mainly  with 
chemical  interactions,  in  which  the  element  either  passes  from  one  state  of  com- 
bination to  another  without  change  of  valence,  or  goes  over  into  a  compound 
of  another  regular  series  with  another  regular  valence.  It  is  not  a  matter  of 
statics.  Hence,  questions  as  to  the  magnitude  of  the  valence  in  isolated  com- 
pounds like  FeaO4,  N2O,  and  so  forth,  are  at  present  of  minor  importance. 

A  definition  of  valence  differing  from  those  given  above  is  preferred  by  many 
chemists.  The  atomic  weight  of  a  univalent  element  can  hold  but  one  unit  of 
another  element  in  combination.  Thus,  the  weight  of  chlorine  represented  by  Cl 
can  hold  but  one  H  or  one  Na  in  combination.  An  atomic  weight  of  a  bivalent 
element,  although  it  combines  with  but  one  unit  of  another  bivalent  element,  may 
hold  as  many  as  two  units  of  a  univalent  element  in  combination.  But  it  cannot 
hold  more.  A  unit  of  a  trivalent  element,  however,  may  hold  as  many  as  three 
units,  provided  the  other  element  is  univalent.  In  this  point  of  view  the  valence 
of  an  element  is  the  maximum  capacity  of  its  atomic  weight  to  hold 
atomic  weights  of  other  elements  in  combination. 

Oddities  Connected  with  Valence.  —  When  the  conception  of  valence 
was  first  evolved,  it  was  taken  for  granted,  quite  mistakenly,  that  each  element 
had  only  on  3  valence.  Hence,  when  pairs  of  compounds  like  CuCl  and  CuCl2,  or 
HgCl  and  HgCl2,  or  FeCl2  and  FeCl3  were  considered,  and  it  was  assumed  that 
copper  was  always  bivalent,  mercury  bivalent,  and  iron  trivalent,  the  formulae 
had  to  be  distorted  to  fit  this  assumption.  Hence  the  formula  of  cuprous  chloride 
was  written,  and  is  still  often  written,  Cu2Cl2. •  The  formulas  of  the  two  chlorides, 
written  graphically  (p.  131)  then  became:  Cl—  Cu  —  Cu—  Gland  Cl—  Cu—  Cl. 
In  both,  each  atom  of  copper  was  holding  (on  paper)  two  other  atoms,  and  was 
therefore  bivalent  (on  paper) .  Similarly  doubling  HgCl  gave  Cl  —  Hg  —  Hg  —  Cl, 

cix  /ci 

and  doubling  FeCl2  gave  Fe  —  Fe         ,  where  the  mercury  appeared  to 

C1X  XC1 

be  bivalent  and  the  iron  trivalent.  Later,  however,  when  elements  like  indium 
were  found  to  have  three  valences,  e.g.,  InCl,  InCl2,  and  InCla,  and  the  com- 
pounds of  manganese  could  not  be  classified  unless  five  different  valences  were 
admitted,  and  many  other  elements  turned  out  to  be  multi-valent,  the  idea  of 
a  fixed  valence  for  each  element  was  given  up.  It  is  entirely  contrary  to  the 
scientific  method  to  invent  or  to  distort  facts  so  as  to  procure  support  for  a 
notion  which  is  a  mere  assumption,  and  not  founded  upon  experiment.  Yet, 
the  history  of  the  science  shows  that  such  errors  were  frequently  committed. 

The  fact  is  that  a  free  element,  like  mercury,  has  no  valence,  for  its  atomic 
weight  is  not  combined  with  even  one  equivalent  of  any  other  element.  When 
combined,  mercury  shows  two  different  valences,  as  in  Hg2O  and  HgO,  HgCl  and 
HgCl2,  HgN03  and  Hg(NO3)2. 

It  is  only  in  the  chemistry  of  carbon  that  the  prejudice  in  favor  of  a  single 
valence  still  persists.  All  chemists  admit  that  in  carbon  monoxide  CO  the 


138  INORGANIC  CHEMISTRY 

carbon  is  bivalent.  But  not  all  chemists  admit  that  it  is  also  bivalent  in  ful- 
minic  acid  H  -  O  —  N  =  C  and  in  the  isonitriles  R  —  N  =  C.  For  the  unsatu- 
rated  compounds,  like  ethylene  C2H4,  and  acetylene  C2H2,  all  chemists  write 
the  formulae  H2  =  C  =  C  =  H2,  and  H  —  C  =  C  —  H,  although  there  is  at 
present  no  experimental  evidence  that  the  formulae  H2  =  C  —  C  =  H2  and 
H— C  —  C—  H,  in  which  the  carbon  is  trivalent  and  bivalent,  respectively, 
do  not  represent  the  facts  equally  well,  so  far  at  least  as  valence  is  concerned. 
When  triphenylmethyl  (C6H6)3C  was  discovered  by  Gomberg,  it  seemed  as  clear 
that  in  this  substance  one  atom  of  carbon  was  trivalent,  as  that  the  copper  in 
Cul  is  univalent;  but  violent  efforts  were  made  to  avoid  this  obvious  conclusion. 
For  example,  it  was  suggested  that  three  of  the  valences  held  one  C6H6  group 
each,  and  that  the  fourth  valence  was  divided  amongst  the  three  groups  (partial 
valence).  But  valence  goes  by  multiples  of  one  equivalent,  and  this  idea  in- 
volved splitting  one  valence  into  thirds  of  an  equivalent.  The  affinity  of  the 
carbon  atoms  is  doubtless  all  divided  amongst  the  three  groups,  so  long  as  no 
fourth  group  is  present  to  share  it,  for  affinity  is  not  divided  into  any  definite 
number  of  units.  This  is  a  case  of  unconsciously  confusing  valence  and  affinity. 
If  we  venture  thus  to  ignore  the  facts,  then  logically  we  must  assume  that  in 
FeCl2  the  iron  is  trivalent,  and  the  third  valence,  in  the  absence  of  a  third  atom 
of  chlorine,  is  being  employed  on  the  other  two.  Clearly  every  element,  includ- 
ing carbon,  is  entitled  to  use  all  the  different  valences  which  are  arithmetically 
possible,  up  to  the  maximum  which  it  ever  exhibits.  The  valence  of  iron  (or 
carbon)  in  a  given  compound  is  not  the  maximum  valence  that  it  shows  in  some 
other  compound,  but  the  valence  it  is  actually  using  in  the  compound  in  question, 
just  as  the  money  a  man  has  in  his  pocket  is  the  amount  actually  there,  and  not 
some  larger  amount  we  think  he  ought  to  have. 

Finally,  it  may  be  noted  that  univalent,  bivalent,  and  trivalent  are  terms 
of  Latin  derivation.  But  tetravalent  is  a  hybrid,  a  mixture  of  Greek  and  Latin. 
One  should  use  either  the  Latin  numerals  or  the  Greek  numerals,  exclusively. 
In  this  book  the  Latin  ones  are  employed  and  the  remaining  valences  are  quadri- 
valent, quinquivalent,  sexivalent,  septivalent,  and  octovalent. 

Inept  Definitions  of  Valence.  —  No  conception  in  chemistry  is  more 
frequently  defined  loosely  or  even  inaccurately  than  is  valence.  Thus,  one  type  of 
definition  runs  somewhat  as  follows:  "The  valence  of  an  element  is  the  number 
of  atoms  it  can  combine  with  or  displace."  Before  criticising  a  general  statement, 
it  is  well  to  substitute  concrete  things  for  the  general  terms.  "The  valence  of 
oxygen  is  the  number  of  atoms  oxygen  can  combine  with  or  displace."  The 
number  of  atoms,  say  of  hydrogen,  with  which  oxygen  can  combine  depends 
entirely  on  how  much  oxygen  is  available.  A  ton  of  oxygen  will  combine  with 
many  more  than  an  ounce,  and  an  ounce  with  many  more  than  an  atom.  But 
this  type  of  definition  omits  to  specify  the  amount  of  oxygen,  namely  the  atomic 
weight,  or  atom,  to  which  the  definition  of  valence  alone  applies.  Oxygen,  as  a 
substance,  is  a  general  term  —  unlimited  as  to  amount  —  so  that  according  to 
this  definition,  the  valence  of  every  element  must  be  infinity!  The  valence  is  not 


HYDROGEN  139 

a  property  of  the  element  in  general,  but 'only  of  its  atomic  weight  (or  atom),  and 
so  the  atomic  weight  must  be  mentioned.  "The  valence  of  an  atom  of  an  element 
is  the  (maximum)  number  of  atoms  it  can  combine  with  or  displace."  Or  "the 
valence  of  an  element  is  the  (maximum)  number  of  atoms  its  atom  can  combine 
with,  etc." 

Almost  every  published  discussion  of  valence  employs  the  word  power. 
Thus:  "Valence  is  the  power  of  the  atom  of  an  element  to  hold  atoms  of  other 
elements  in  combination."  Now  power  is  defined  in  physics  as  the  rate  at  which 
a  machine  does  work.  Valence  —  say  that  in  a  pound  of  salt  —  does  no  work 
whatever.  Hence,  when  the  word  power  is  used,  the  valence  of  every  element 
becomes,  ipso  facto,  zero!  This  definition  goes  to  the  opposite  extreme  of  error 
from  the  last.  The  word  suggests,  and  is  perhaps  intended  to  suggest,  that  va- 
lence is  a  measure  of  the  force  with  which  the  atoms  are  held  together  in  a  com- 
pound. But  valence  has  nothing  to  do  with  the  force  —  which  is  a  question  of 
affinity  —  but  solely  with  the  quantity  of  matter  held  in  combination.  Thus, 
gold  is  trivalent  (AuCls),  but  the  chloride  can  be  decomposed  by  gentle  heating, 
while  sodium  is  univalent,  yet,  its  chloride  NaCl  can  be  vaporized  at  a  high  tem- 
perature without  decomposition.  So  an  atom  of  gold  holds  three  times  as  much 
chlorine  in  combination  as  does  an  atom  of  sodium  (and  is  trivalent),  although 
it  holds  it  very  feebly,  or,  as  one  might  say,  with  little  force  or  "power."  Valence 
is  the  number  of  equivalents  combined  with  the  given  atomic  weight :  it  is  simply 
a  question  of  capacity  (see  definition,  p.  137).  A  tank  might  have  a  capacity  of  a 
hundred  million  gallons  of  water,  yet,  if  the  tank  were  close  to  the  sea  level,  the 
available  power  would  be  almost  zero.  Capacity  and  force  and  power  are  entirely 
different  conceptions.  The  atom  of  some  element  might  be  able  to  hold  eight 
atoms  of  hydrogen  in  combination  —  that  would  be  its  capacity  or  valence  — 
yet,  it  might  hold  them  so  feebly  (with  so  little  force  or  "power")  that  the  com- 
pound would  give  off  hydrogen  at  room  temperature,  and  would  have  to  be  kept 
in  a  refrigerator.  Correct  use  of  technical  terms  is  indispensable  in  chemistry, 
and  capacity,  not  power,  is  the  term  applicable  to  valence. 

Exercises.  —  1.  Make  equations  for  reactions  in  which  hydro- 
gen is  liberated  by  the  action  of:  (a)  hydrochloric  acid  and  magne- 
sium, giving  MgCl2,  (6)  steam  and  zinc,  giving  ZnO. 

2.  Make  an  equation  for  the  action  of  heat  on  manganese  dioxide 
MnO2,  giving  oxygen  and  Mn3O4. 

3.  What  must  be  the  relative  rates  of  diffusion  of  hydrogen  and 
of  carbon  dioxide? 

4.  Make  equations  to  represent,  (a)  the  reduction  of  lead  dioxide 
PbC>2  by  hydrogen;    (b)  the  actions  of  aluminium  upon  cold  water 
and  upon  steam  at  a  red  heat. 

5.  Which  are  the  components  (p.  7)  of  dilute  sulphuric  acid 
and  which  are  the  constituents  of  sulphuric  acid? 

6.  What  are  the  valences  of  the  negative  radicals  of  phosphoric 


140  INORGANIC  CHEMISTRY 

acid  H3PO4,  and  of  acetic  acid  (p.  117)?  What  must  be  the  formulae 
of  calcium  phosphate,  cupric  acetate  (p.  118),  aluminium  phosphate, 
ferrous  carbonate,  ferrous  sulphate,  cupric  chloride? 

7.  What  is  the  valence  of  phosphorus  in  phosphoric  anhydride 
(p.  90)?     What  must  be  the  formulae  of,  (a)  the  corresponding  chlo- 
ride and  the  sulphide  of  phosphorus,  and  (6)  of  aluminium  oxide? 

8.  What  are  the  valences  of  the  elements  in  the  following:  LiH, 
NH3,  SeH2,  BN? 

9.  What  are  the  valences  of  the  metals  and  radicals  in  the  follow- 
ing:  Pb(NO3)2,  Ce(SO4)2,  KC1,  KMnO4  (potassium  permanganate)? 
Name  all  the  substances  in  8  and  9. 

10.  Write  the  formulae  of  ferrous  and  ferric  oxides,  of  ferrous 
and  ferric  nitrates,  of  stannous  and  stannic  sulphides. 

11.  One  gram  of  a  quadrivalent  element  unites  with  0.27  g.  of 
oxygen.     What  is  the  atomic  weight  of  the  element? 


CHAPTER  VIII 
WATER 

THE  great  quantity  of  water  which  occurs  in  nature  makes  it  one 
of  the  most  familiar  chemical  substances.  The  ocean  covers  about 
three-fourths  of  the  surface  of  the  earth,  and  in  most  habitable  regions 
lakes  and  streams  abound.  The  "dry  "  land  is,  fortunately,  far  from 
being  really  dry,  Water  is  found  also  in  the  bodies  of  both  animals 
and  plants  in  large  quantities,  and  is  indeed  essential  to  the  working 
of  living  organisms. 

Natural  Waters.  —  The  water  found  in  nature  varies  greatly  in 
the  amount  of  foreign  material  which  it  contains.  Sea-water  holds 
about  3.6  per  cent  of  solid  matter  in  solution,  while  rain-water  is  the 
purest  natural  water.  Even  rain-water  contains  foreign  matter,  how- 
ever. When  we  heat  it,  bubbles  of  gas  form  on  the  sides  of  the  vessel, 
showing  that  oxygen,  nitrogen,  and  carbon  dioxide  from  the  air  have 
been  dissolved  by  the  water  as  it  fell.  On  evaporating  a  considerable 
mass  of  such  water,  we  find  that,  aside  from  dust,  crystals  of  chemical 
substances,  such  as  ammonium  nitrate,  may  be  recognized  in  the 
residue.  Of  well  and  surface  waters,  some  which  contain  calcium 
sulphate,  calcium  bicarbonate,  and  compounds  of  magnesium  in 
solution  are  described  as  hard.  Others  contain  compounds  of  iron, 
and  still  others  are  effervescent  and  give  off  carbon  dioxide.  These 
are  called  mineral  waters.  All  of  the  dissolved  substances  are 
obtained  by  the  water  in  its  progress  over  or  under  the  surface  of  the 
ground. 

Water  which  is  to  be  used  for  domestic  purposes  is  examined,  not 
only  to  ascertain  the  amount  of  the  ingredients  which  produce  hard- 
ness, but  also  with  reference  to  the  proportion  of  organic  matter  which 
it  may  hold  in  solution.  This  usually  gains  access  to  the  water  by 
admixture  of  sewage  (p.  92).  It  is  not  the  organic  matter  itself  which 
is  deleterious,  but  the  bacteria  of  putrefaction  and  disease  which  are 
likely  to  accompany  it.  If  the  water  contains  clay  in  suspension,  the 
bacteria  are  largely  attached  to  the  particles  of  clay,  but  organic 
matter  and  bacteria  may  be  present  in  water  which  looks  perfectly 


142 


INORGANIC   CHEMISTRY 


clear.     Inoculation  of  culture  media  with  the  water  can  alone  show 
whether  or  not  bacteria  are  present. 

The  foreign  materials  which  water  may  contain  are  divisible  into 
two  kinds  —  dissolved  matter  and  suspended  matter.  No  natural 
water  is  entirely  free  from  either  of  these  varieties  of  impurity. 

Purification  from  Suspended  Matter  and  Bacteria.  —  The 

suspended  impurities  may  be  removed  by  filtration.  On  a  large 
scale,  beds  of  gravel  are  employed,  but  this  treatment  will  not  remove 
all  bacteria.  In  many  cases  small  amounts  of  alum,  or  alum  and 
lime,  or  ferrous  sulphate  (copperas)  and  lime,  are  added.  These 
produce  slimy  precipitates  which  assist  in  the  elimination  of  fine, 
suspended  inorganic  and  organic  matter,  including  practically  all 
bacteria.  This  is  called  the  coagulation  treatment  (q.v.).  The  whole 
suspended  matter  is  then  allowed  to  settle,  which  it  does  very  quickly, 
in  large  reservoirs.  The  remaining  organisms  may  be  destroyed  by 
adding  a  little  bleaching  powder  (q.v.),  or  chlorine-water,  before  the 
water  is  distributed.  Ozone  (q.v.)  and  ultra-violet  light 
are  used  for  the  same  purpose,  the  latter,  for  example,  at 
Rouen.  The  ultra-violet  light,  which  is  a  light  of  very 
short  wave  length,  is  produced  by  an  electric  arc  passing 
through  mercury  vapor  in  a  container  of  quartz  (glass 
absorbs  and  destroys  this  light).  The  water  must  first 
be  filtered  so  that  the  light  may  be  able  to  penetrate  it. 
A  3-ampere  lamp  on  a  220-volt  circuit  will  kill  the  colon 
bacilli,  organisms  associated  with  those  which  produce 
typhoid  fever,  in  a  layer  of  water  10  cm.  thick  in  1  second 
and  in  a  layer  40  cm.  thick  in  fifteen  seconds.  The  flow 
of  the  water  is  regulated  so  as  to  permit  sufficiently  long 
exposure  to  the  rays. 

In  the  household,  the  Pasteur  filter  is  the  most  com- 
pact and  efficient  appliance.  The  water  enters  at  the 
top  (Fig.  52)  and  is  forced  inwards  by  its  own  pressure 
through  the  pores  of  a  cylinder  of  unglazed  porcelain. 
The  cylinder  must  be  taken  out,  and  its  exterior  cleaned 
daily  with  a  brush,  to  remove  the  mud  and  organisms 
which  collect  on  its  outer  surface.  If  this  is  not  done,  the  organisms 
multiply  and  soon  the  filter  pollutes  the  water  instead  of  purify- 
ing it. 

Most  organisms  can  be  killed  by  boiling  unfiltered  water  for  10  or 
15  minutes,  although  a  second  boiling  is  needed  in  some  cases. 


FIG.  52. 


WATER  143 

Purification  from  Dissolved  Matter.  —  Filtration  does  not 
remove  dissolved  matter,  and  therefore  does  not  soften  hard  water 

(**)> 

Pure  water  for  chemical  purposes  is  prepared  by  distillation  and, 
in  fact,  liquids  other  than  water  are  usually  purified  by  the  same 
process  (Fig.  20,  p.  43).  The  steam  is  condensed  by  cold  water 
circulating  in  the  jacket,  and  contains  at  first  only  gases  dissolved 
from  the  air.  The  dissolved  solids  remain  in  the  flask.  Distilled 
water  quickly  dissolves  traces  of  glass  or  porcelain,  so  that  the  purest 
water  is  obtained  by  using  quartz  or  platinum  for  the  condenser 
tube  and  receiving  vessel.  Tin  is  the  best  of  the  less  expensive 
materials. 

That  glass  dissolves  in  water  is  easily  shown  by  shaking  some 
pulverized  glass  with  distilled  water  for  a  few  seconds  and  adding  a 
drop  of  phenolphthalein  solution  (see  Indicators).  The  alkaline 
reaction  of  the  dissolved  sodium  silicate  gives  rise  to  a  strong  pink 
color  [Lect.  exp.]. 

Physical  Properties  of  Water.  —  When  we  view  a  white  object 
through  a  deep  layer  of  water  we  find  that  the  liquid  has  a  blue  or 
greenish-blue  color.  At  a  pressure  of  760  mm.,  water  exists  as  a 
liquid  between  0°  and  100°.  Below  0°  it  becomes  solid,  above  100° 
a  gas.  Of  all  chemical  substances  it  is  the  one  which  we  use  most,  so 
that  familiarity  with  its  physical  properties,  discussed  below,  is  indis- 
pensable to  the  chemist.  It  will  serve  also  as  a  typical  liquid,  since 
it  differs  from  others  only  in  details. 

The  quantity  of  heat  required  to  raise  one  gram  of  water  one 
degree  in  temperature,  at  15°,  is  called  a  calorie,  and  is  the  unit  of 
.heat.  The  specific  heat  of  any  substance  being  also  the  quantity  of 
heat  required  to  raise  the  temperature  of  one  gram  of  the  substance 
one  degree,  the  specific  heat  of  water  is  1.  The  values  for  other 
substances  are  all  smaller  (e.g.,  limestone  0-2).  Thus  the  tempera- 
ture of  large  masses  of  water,  such  as  lakes  and  seas,  changes  more 
slowly,  and  within  a  smaller  range,  than  that  of  the  rocks  and  soil 
composing  the  land.  The  more  constant  temperature,  of  the  water 
tends  to  regulate  that  of  the  air,  and  hence  the  climate  of  islands  is 
less  variable  from  season  to  season  than  is  that  of  a  continent. 

The  weight  of  a  cubic  centimeter  of  water  at  4°  gives  us  our  unit,  the  gram.  A 
kilogram  of  water  at  0°  occupies  1.00013  liters,  or  0.13  c.c.  more  than  at  4°  C.  A 
kilogram  of  ice  at  0°  occupies  1  09083  liters,  or  90.7  c.c.  more  than  an  equal  weight 
of  water  at  0°.  The  volume  of  one  kilogram  of  water  at  100°  is  1.0432  liters. 


144  INORGANIC  CHEMISTRY 

Ice.  —  The  raising  or  lowering  of  the  temperature  of  a  gram  of 
water  through  one  degree  corresponds  to  the  addition  or  removal  of 
one  calorie  of  heat.  The  conversion,  however,  of  a  gram  of  water  at 
0°  to  a  gram  of  ice  at  0°  requires  the  removal  of  79  calories  of  heat. 
The  mere  melting  of  a  gram  of  ice  causes  an  absorption  of  heat  to  the 
same  amount.  This  is  called  the  heat  of  fusion  of  ice.  At  0°  a  mix- 
ture of  ice  and  water  will  remain  in  unchanged  proportions  indefinitely. 
Any  cause  which  tends  permanently  to  lower  or  raise  the  temperature 
by  a  fraction  of  a  degree,  however,  will  bring  about  the  disappearance 
of  the  water  or  of  the  ice,  respectively.  This  temperature  (0°)  is 
called  the  melting-  or  the  freezing-point  of  water.  Properties  of  this 
kind,  marked  by  transition  points  from  one  state  to  another,  are  much 
used  in  chemistry  for  keeping  other  bodies  or  systems  at  a  constant 
temperature  during  measurement  or  observation.  A  mixture  of  ice 
and  water  surrounding  a  body,  when  kept  in  constant  agitation,  will 
automatically  maintain  the  body  at  a  fixed  temperature  (0°)  so  long 
as  both  components  hold  out. 

Water  can  easily  be  cooled  below  0°  (supercooled)  without  begin- 
ning to  freeze,  unless  it  is  stirred  or  " inoculated"  by  the  addition  of 
a  piece  of  ice.  Hence,  the  freezing-point  is  not  defined  as  the  point 
at  which  ice  begins  to  form,  for  that  point  varies,  and  is  always 
below  0°,  but  as  the  temperature  of  a  well-stirred  mixture  of  ice 
and  water. 

Steam.  —  At  atmospheric  pressure,  water  passes  into  steam 
rapidly  at  100°,  but  at  lower  temperatures,  and  even  when  frozen,  it 
does  the  same  thing  more  slowly.  It  changes  into  steam,  however, 
only  when  the  necessary  supply  of  heat  is  forthcoming.  One  gram  of 
water  at  100°,  for  example,  in  turning  into  a  gram  of  steam  at  100°,  • 
takes  up  540  calories.  This  is  called  its  heat  of  vaporization.  Steam, 
in  fact,  contains  much  more  internal  energy  than  an  equal  weight  of 
water  at  the  same  temperature,  just  as  water,  in  turn,  contains  more 
energy  than  ice. 

Steam  is  a  colorless,  invisible  gas.  The  visible  cloud  of  fog,  seen 
when  steam  escapes  into  cold  air,  is  composed  of  minute  drops  of 
water,  formed  by  condensation,  and  visible  because  they  have 
surfaces  which  reflect  light. 

The  temperature  of  100°  is,  like  the  melting-point  of  ice,  an  im- 
portant transition  point.  It  is  less  exactly  recoverable  by  simply 
keeping  a  vessel  full  of  water  in  ebullition,  however,  because  the 
natural,  and  often  considerable,  variations  in  the  pressure  of  the 


WATER 


145 


atmosphere  affect  it  more  markedly  than  they  do  the  melting-point 
of  ice.  Near  to  100°,  the  boiling-point  rises  or  falls  about  0.037°  for 
1  mm.  change  in  pressure.  On  the  top  of  Mont  Blanc  water  boils  at 
84°.  The  melting-point  of  ice  is  lowered  only  0.0075°  by  an  increase 
in  pressure  from  760  mm.  to  2  atmospheres  (1520  mm.) 

Most  substances  are  known  in  three  different  states  of  aggrega- 
tion, solid,  liquid,  and  gaseous.  There  is  no  magic  about  the  num- 
ber, three,  however.  Thus,  sulphur  has  a  vapor  state,  two  liquid 
states,  and  at  least  four  different  solid  forms.  There  are  even  five 
different  forms  of  ice,  and  most  solids  probably  exist  in  several 
different  states. 

Vapor  Pressure  and  Aqueous  Tension.  —  The  quantity  of 
vapor  given,  off  by  a  substance  is  defined  by  the  gaseous  pressure  it 
exercises.  This  is  called  the  vapor  pressure  of  the  substance. 

The  most  significant  fact  about  vapor  pressure  is  that,  when  ex- 
cess of  the  liquid  is  present,  the  pressure  of  the  vapor  quickly  reaches 
a  definite  maximum  value  for  each  temperature.  In  the  absence 
of  excess  of  the  water,  less  than  this  maximum  pressure 
may  exist.  More  than  the  maximum  pressure  proper 
to  a  given  temperature,  if  produced  by  compression, 
cannot  be  maintained,  however,  for  a  part  of  the 
vapor  condenses  to  the  liquid  state.  The  magnitude 
of  this  maximum  vapor  pressure,  at  a  given  temper- 
ature, depends  on  the  ability  of  the  particular  liquid 
to  generate  vapor.  This  maximum  vapor  pressure  is 
held,  therefore,  to  represent  the  vapor  tension  of  the 
liquid,  at  the  given  temperature,  and  this  is  a  specific 
property  of  the  substance. 

The  vapor  tension  may  be  shown  by  allowing  a  few 
drops  of  water  to  ascend  into  the  vacuum  at  the  top  of 
a  barometric  column  (Fig.  53).  The  tube  on  the  left 
shows  the  mercury  when  nothing  presses  on  its  surface. 
The  tube  on  the  right  shows  the  result  of  admitting 
the  water.  The  pressure  of  the  atmosphere  being  the 
same  for  both,  the  smaller  height  of  mercury  which  now  suffices  to 
counterbalance  it  shows  that  something,  which  can  be  nothing  but 
the  water  vapor,  is  pressing  on  the  surface  of  the  mercury,  and  makes 
up  the  rest  of  the  total  stress  needed.  The  difference  in  the  height 
of  the  two  columns  gives  the  value  of  this  pressure,  which  we  calJ 
the  vapor  pressure,  of  the  water. 


FIG.  53. 


146  INORGANIC  CHEMISTRY 

With  excess  of  liquid  water,  the  value  is  that  of  the  vapor  tension 
of  the  liquid,  called,  in  the  case  of  water,  the  aqueous  tension. 

The  jacket  surrounding  the  tube  on  the  right  enables  us,  by  adding 
ice  or  warm  water,  to  keep  the  water  that  is  admitted  to  the  vacuum, 
and  the  parts  of  the  apparatus  immediately  in  contact  with  it,  at  any 
temperature  between  0°  and  100°.  When  ice  is  used  outside,  and  a 
piece  of  it  is  introduced  into  the  vacuum,  the  vapor  it  gives  off  quickly 
reaches  a  pressure  of  4.5  mm.  The  vapor  pressure  of  the  ice  takes  the 
place  of  4.5  mm.  of  mercury  in  balancing  the  atmospheric  pressure, 
and  so  the  mercury  column  falls  by  this  amount.  Similarly,  water  at 
10°  causes  a  fall  of  9.1  mm.  and  at  20°  of  17.4  mm.,  so  that  these 
represent  the  mercury-height  values  of  the  aqueous  tension  at  these 
temperatures.  With  ether,  instead  of  water,  at  10°,  the  fall  is  28.7 
mm.  The  quantity  of  water  used  makes  no  difference,  so  long  as  a 
little  more  is  present  than  is  required  to  fill  the  available  space  with 
vapor.  Of  course,  if  a  large  amount  is  admitted,  its  dead  weight  will 
take  the  place  of  an  equal  weight  of  mercury  in  balancing  the  pressure 
of  the  air.  If  there  is  a  measurable  column  of  water,  its  height  must 
be  divided  by  13.6  (the  density  of  mercury),  and  counted  as  if  it  were 
part  of  the  mercury. 

With  water  at  higher  temperatures  the  fall  of  the  mercury  column 
becomes  much  greater.  At  50°  it  is  92  mm.,  at  70°  it  is  233.3  mm.,  at 
90°  it  is  525.5  mm.,  and  at  100°  it  is  760  mm.  At  the  boiling-point, 
therefore,  the  aqueous  tension  takes  the  place  of  the  whole  barometric 
column,  and  is  equal  to  the  average  air  pressure.  At  121° 
the  aqueous  tension  is  two  atmospheres,  at  180°  it  is  ten 
atmospheres  (see  Appendix  IV). 

There  is  another  standpoint  from  which  these  phe- 
^  nomena  may  be  viewed.  Water  vapor  can  exist  at  10° 
only  if  the  pressure  upon  it  is  9.1  mm.  or  less.  If  we 
imagine  the  water  placed  in  a  cylinder  closed  by  a  fric- 
tionless,  weightless  piston  (Fig.  54),  then  at  10°  the  piston 
will  remain  at  rest  whether  we  place  it  high  or  low,  pro- 
vided it  is  loaded  with  a  weight  exactly  equal  to  that  of 
a  layer  of  mercury  9.1  mm.  thick  covering  its  whole  area. 
We  speak  of  such  a  system  as  being  in  equilibrium  (see  Chap.  IX). 
With  a  less  weight  the  piston  will  move  slowly  upwards,  as  the  vapor 
continually  given  off  by  the  water  presses  upon  it,  until  it  reaches  the 
top  or  the  water  all  evaporates.  Conversely,  if  it  bears  a  greater  load, 
it  will  move  down  and  the  vapor  will  condense  on  the  walls  and 
bottom  of  the  cylinder  until  the  piston  comes  in  contact  with  the 


FIG.  54. 


WATER  147 

water  itself  and  the  vapor  is  all  abolished.  These  conceptions  will 
find  constant  application  not  only  to  physical  but  also  to  chemical 
phenomena.  The  expression: 

Aq  (liq.)<=±Aq(vap.) 

is  used  to  represent  the  state  of  equilibrium  in  a  system  like  the  above. 
When  water  at  a  certain  temperature  has  given  the  full  amount  of 
water  vapor  to  the  space  above  it  that  its  aqueous  tension  permits,  we 
say  that  the  space  is  saturated  with  vapor.  That  concentration  of 
vapor  which  constitutes  saturation  varies  with  the  temperature  of  the 
water  and  depends  therefore  solely  on  the  ability  of  the  water  to  give 
off  vapor.  It  has  nothing  to  do  with  the  size  of  the  space,  and  is  even 
independent  of  other  gases  the  space  may  already  contain.  Thus,  if 
a  little  air  is  first  placed  above  the  dry  mercury  (Fig.  53),  causing  it 
to  fall,  the  additional  depression  produced  by  adding  water  is  the 
same  as  if  the  air  had  been  absent  (p.  111). 

Water  Vapor  in  the  Air.  —  The  space  immediately  above  the 
surface  of  the  ground,  which  is  mainly  occupied  by  atmospheric  air, 
is,  on  an  average,  less  than  two-thirds  saturated  with  water  vapor. 
That  is  to  say,  such  air,  when  enclosed  in  a  vessel  containing  water, 
will  take  up  about  one-half  more  than  it -already  contains.  The 
vapor  of  water  at  100°  in  an  open  vessel  displaces  the  air  entirely  and, 
if  the  required  heat  of  vaporization  is  furnished,  the  liquid  boils. 

The  water  present  in  the  air  plays  an  important  part  in  many 
chemical  phenomena,  as  we  shall  see.  All  our  substances  and  appa- 
ratus have  traces  of  water  condensed  on  their  surfaces.  This  water  is, 
in  a  sense,  in  an  abnormal  condition,  for  it  does  not  evaporate  even  in 
dry  air.  It  is  observed  to  pass  off  in  vapor,  however,  when  we  have 
occasion  to  heat  the  substance  or  apparatus. 

Water  as  a  Solvent.  —  One  of  those  physical  properties  of  water 
which  are  most  used  in  chemical  work  is  its  tendency  to  dissolve  many 
substances.  This  subject  is  so  important  and  extensive  that  we  shall 
presently  devote  a  complete  chapter  to  some  of  its  simpler  and  more 
familiar  aspects. 

Chemical  Properties  of  Water.  —  Water  is  so  very  frequently 
used  in  chemical  experiments  in  which  it  is  a  mere  mechanical  ad- 
junct, that  the  beginner  has  difficulty  in  distinguishing  the  cases  in 
which  it  has  itself  taken  part  in  the  chemical  interaction.  The  rather 


148  INORGANIC  CHEMISTRY 

limited  list  of  kinds  of  chemical  activity  it  can  show  should  therefore 
receive  careful  notice: 

1.  Water  is  a  relatively  stable  substance. 

2.  It  combines  with  many  oxides,  forming  bases  or  acids. 

3.  It  combines  with  many  substances,   chiefly  salts,   forming 
hydrates. 

4.  It  interacts  with   some  substances  in  a  way  described  as 
hydrolysis.     This  property  will  not  be  discussed  until  a  characteristic 
example  is  encountered. 

Perhaps  we  should  add  that  steam  at  a  high  temperature  oxidizes 
elements  which  readily  combine  with  oxygen.  For  example,  it  turns 
iron  into  the  magnetic  oxide  (p.  116).  At  such  high  temperatures, 
however,  the  water  is  partially  resolved  into  a  mixture  of  hydrogen  and 
oxygen  and,  the  latter  being  the  more  active  of  the  two  elements, 
towards  iron,  the  oxidizing  effects  predominate.  Even  other  com- 
pounds containing  oxygen  will  give  exactly  the  same  results.  Hence 
this  cannot  be  regarded  as  a  property  of  water  itself. 

Water  a  Stable  Compound:  Dissociation.  —  In  the  case  of 
a  compound,  the  first  chemical  property  to  be  given  is  always,  whether 
the  substance  is  stable  or  unstable.  Usually,  the  specification  is  in 
terms  of  the  temperature  required  to  produce  noticeable  decomposi- 
tion. Thus,  potassium  chlorate  gives  off  oxygen  at  a  low  red  heat. 
Now,  water  vapor,  when  heated,  is  partially  decomposed  into  hydro- 
gen and  oxygen,  yet,  at  1882°  only  1.18  per  cent  is  broken  up,  and  at 
2000°  the  decomposition  reaches  only  1.8  per  cent.  When  the  tem- 
perature is  lowered,  the  gases  recombine  to  form  water.  Two  arrows 
in  the  equation  indicate  that  the  action  may  proceed  in  either  direc- 
tion —  is  reversible: 


A  decomposition  which  increases  when  the  temperature  is  shifted 
in  one  direction  (usually,  though  not  always,  upward),  and  reverses 
itself  by  recombination  of  the  constituents  when  the  temperature  is 
displaced  in  the  other  direction  (usually  downward),  is  called  a 
dissociation  or  a  thermal  dissociation.  The  decomposition  of  potas- 
sium chlorate  (p.  83)  is  not  a  dissociation,  because  it  is  not  reversible: 
oxygen  gas  will  not,  under  any  known  circumstances,  unite  directly 
with  potassium  chloride. 

The  word  unstable  is  frequently  misused.  Thus,  calcium  chloride  is 
sometimes  said  to  be  unstable,  because,  when,  exposed  to  the  air,  it  takes  up 


WATER  149 

moisture  and  finally  dissolves  in  the  water  thus  acquired.  Or  phosphorus 
trichloride  is  said  to  be  unstable  because  it  interacts  with  the  moisture  in  the  air, 
or  ferrous  sulphate  is  said  to  be  unstable  because  it  is  oxidized  by  the  oxygen 
in  the  air.  Almost  every  substance,  however,  will  undergo  change  readily,  if  a 
properly  selected  substance  is  allowed  to  come  in  contact  with  it.  Quicklime,  for 
example,  is  quickly  slaked  by  water.  According  to  this  erroneous  view  prac- 
tically every  substance  is  unstable.  Quicklime  can  be  distilled,  without  decom- 
position, at  the  temperature  of  the  electric  arc,  and  is  extremely  stable.  Stability 
refers  solely  to  the  effect  of  energy,  usually  in  the  form  of  heat,  upon  the  substance, 
while  the  latter  is  carefully  isolated  from  all  other  substances.  The  interactions 
of  the  substance  with  the  components  of  the  air,  and  with  other  reagents,  must 
be  recorded  as  separate  chemical  properties. 

Union  of  Water  with  Oxides.  —  1.  Sodium  oxide  Na^O  unites 
violently  with  water  to  form  sodium  hydroxide: 

Na20  +  H2O-»2NaOH. 

The  slaking  of  quicklime  is  a  more  familiar  action  of  the  same  kind: 
CaO  +  H20  ->  Ca(OH)2. 

No  other  products  are  formed.  The  clouds  of  steam  produced  in  the 
second  instance  are  due  to  evaporation  of  a  part  of  the  water  by  the 
heat  produced  in  the  formation  of  calcium  hydroxide.  The  aqueous 
solutions  of  these  two  products  have  a  soapy  feeling,  and  turn  red  lit- 
mus blue  (see  Indicators),  and  the  substances  therefore  belong  to  the 
class  of  alkalies  or  bases.  Very  many  hydroxides  which  are  of 
the  same  nature,  for  example  ferric  hydroxide  Fe(OH)3  and  tin  hy- 
droxide Sn(OH)2,  are  formed  so  slowly  by  direct  union  of  the  oxide 
and  water  that  they  are  always  prepared  in  other  ways.  The  oxides 
which,  with  water,  form  bases  are  called  basic  oxides. 

2.  Some  oxides,  although  they  unite  with  water,  give  acids,  which 
are  products  of  an  entirely  different  character.  Phosphorus  pent- 
oxide  P2O6  and  sulphur  dioxide  SO2  (p.  88)  are  of  this  class  and  yield 
phosphoric  acid  and  sulphurous  acid,  respectively: 


S02  +H20   <=»H2S03. 

If  the  product  is  not  volatile,  it  may  be  obtained  by  evaporating  the 
excess  of  water.  In  the  case  of  sulphurous  acid,  the  above  action  is 
reversed  by  evaporation  and  the  sulphur  dioxide  and  water  both  pass 
off;  in  that  of  phosphoric  acid,  the  white  crystalline  acid  is  obtained. 
In  consequence  of  their  relation  to  the  acid,  differing  from  it  in  not 


150  INORGANIC  CHEMISTRY 

containing  the  elements  of  water,  these  oxides  are  often  called  the 
anhydrides  (Gk.  dv,  not,  and  vSa>p,  water),  of  their  respective  acids. 
They  are  also  called  acidic  oxides.  The  acids  are  sour  in  taste  and 
turn  blue  litmus  red. 

These  two  classes  of  final  products  are  so  different  that  we  make 
the  distinction  the  basis  of  classification  of  the  elements  present  in  the 
original  oxides.  The  elements,  like  sodium  and  iron,  whose  oxides 
give  bases,  are  called  metallic  elements;  those,  like  phosphorus, 
whose  oxides  give  acids,  are  called  non-metallic  elements.  The  dis- 
tinguishing words  are  selected  because  the  division  corresponds,  in  a 
general  way  at  least,  with  the  separation  into  two  sets  to  which  merely 
physical  examination  of  the  elementary  substances  would  lead. 

Formerly  the  hydroxides  of  metals  were  termed  "  hydrates,"  and 
the  word  is  still  used  familiarly  by  chemists  in  a  few  cases,  such  as 
potassium  " hydrate"  KOH  and  sodium  " hydrate"  NaOH.  These 
substances,  however,  have  nothing  in  common  with  the  compounds 
properly  known  as  hydrates,  whose  nature  is  discussed  in  the  next 
section. 

Hydrates.  —  Many  substances  when  dissolved  in  water  and  re- 
covered by  spontaneous  evaporation  of  the  solvent  are  found  to  have 
entered  into  combination  with  the  liquid.  The  products,  which  are 
solids,  are  called  hydrates.  That  they  are  regular  chemical  com- 
pounds is  shown  by  the  following  two  facts:  (1)  These  compounds 
show  definite  chemical  composition  expressible  by  formulae  in  terms 
of  chemical  unit  weights  (atomic  weights)  of  the  constituent  elements. 
The  proportions  in  solutions,  and  other  physical  aggregates,  except 
by  chance,  cannot  be  expressed  by  means  of  formulae.  (2)  A  hydrate 
has  physical  properties  entirely  different  from  those  of  the  water  (or 
ice)  and  of  the  other  substance  used  in  preparing  it.  It  is  a  typical 
compound,  formed  by  the  first  variety  of  chemical  change  (p.  12). 
Thus,  cupric  sulphate,  often  called  anhydrous  cupric  sulphate  to 
distinguish  it  from  the  compound  with  water,  is  a  white  substance 
crystallizing  in  shining,  colorless,  needle-like  prisms.  The  penta- 
hydrate  (blue-stone  or  blue  vitriol)  is  blue  in  color,  and  forms  larger 
but  much  less  symmetrical  (asymmetric  or  triclinic)  crystals  (Fig. 
55): 

CuSO4  +  5H20  ->  CuSO4,5H20. 

Often  much  heat  is  given  out  in  the  formation  of  a  hydrate. 
Thus,  in  the  case  of  washing  soda,  the  decahydrate  of  sodium  car- 


WATER  151 

bonate  Na2CO3,10H2O,  the  heat  of  union  (p.  99)  is  +8800  cal. 
This,  however,  does  not  in  itself  show  that  a  chemical  change  has 
occurred.  All  chemical  changes  are  accompanied  by  liberation  or 
absorption  of  heat,  but  physical  changes,  like 
the  condensation  of  steam  and  the  freezing  of 
water,  and  in  many  cases  the  dissolving  of  one 
substance  in  another,  also  involve  the  libera- 
tion of  much  heat. 

The  chemical  properties  show  hydrates  to 
be  relatively  unstable.  When  heated,  the  hy- 
drates, as  a  rule,  lose  none  of  the  constituents 
of  the  original  compound,  but  only  those  of 
the  water  in  the  form  of  water  vapor.  When 
melted,  or  when  dissolved  in  water  the  hydrates  are  partly  dissociated 
into  water  and  the  original  substance.  The  aqueous  solutions 
made  from  the  anhydrous  substances  and  from  the  hydrates  have 
identical  physical  and  chemical  properties.  Hence  the  cheaper  of 
the  two  forms  is  generally  purchased,  and  many  of  the  chemicals  used 
in  laboratories  are  in  the  form  of  hydrates. 

Since  hydrates,  when  they  decompose,  usually  give  up  water,  we 
write  their  formulae,  e.g.,  CuSO4,5H20,  so  that  the  water  and  the 
original  substance  are  separated  by  a  comma.  A  formula  thus  modi- 
fied, so  as  to  show  some  favorite  mode  of  chemical  behavior  of  the 
substance,  is  called  a  reaction  formula.  The  formula  HioCuSOg, 
which  would  show  the  same  proportions  by  weight,  is  never  employed, 
because  its  use  would  disguise  the  relation  of  the  substance  to  cupric 
sulphate. 

The  Dissociation  of  Hydrates:  Efflorescence.  —  The  less 
stable  hydrates  dissociate  very  readily.  Thus,  the  decahydrate  of 
sodium  sulphate,  Na2S04,10H2O  (Glauber's  salt),  loses  all  the  water 
it  contains  (effloresces)  when  simply  kept  in  an  open  vessel.  When 
kept  in  a  closed  bottle,  a  very  little  of  it  loses  water,  and  then  the 
decomposition  ceases.  The  cause  of  this  we  discover  when  a  crystal 
of  the  hydrate  is  placed  above  mercury,  like  the  ice  or  water  in  Fig.  53 
(p.  145).  .It  shows  an  aqueous  tension  which  we  can  measure.  At  9° 
the  value  of  this  is  5.5  mm.  As  its  temperature  is  raised,  the  tension 
increases.  When  the  temperature  is  lowered,  on  the  other  hand,  the 
tension  diminishes,  the  mercury  rises,  and  a  part  of  the  water  enters 
into  combination  again.  Different  hydrates  show  different  aqueous 
tensions  at  the  same  temperature.  For  example,  at  30°,  that  of  water 


152  INORGANIC  CHEMISTRY 

itself  is  31.5  mm.;  strontium  chloride  SrCl2,6H20,  11.5  mm.;  cupric 
sulphate  CuSO4,5H2O,  12.5  mm.;  barium  chloride  BaCl2,2H2O, 
4  mm. 

In  view  of  these  facts,  we  perceive  that'  loss  of  water  by  efflores- 
cence is  like  evaporation,  excepting  that  it  is  a  chemical  decomposition 
and  not  a  physical  process.  Those  hydrates  which,  like  Glauber's 
salt  and  washing  soda  Na2CO3,10H2O,  have  a  vapor  tension  approach- 
ing that  of  water  itself,  lose  their  water  at  ordinary  temperatures  at 
a  rapid  pace.  In  this  connection  we  have  to  remember  that  atmos- 
pheric air  is  always  less  than  two-thirds  saturated  with  water  vapor, 
and  the  partial  pressure  of  this  vapor  opposes  the  dissociation,  and 
tends  to  prevent  the  liberation  of  the  water.  Thus  at  9°,  the  vapor 
tension  of  water  being  8.6  mm.,  the  average  vapor  pressure  of  water 
in  the  atmosphere  will  be  about  5  mm.  Any  hydrate  with  a  greater 
aqueous  tension  than  5  mm.,  at  9°,  such  as  Glauber's  salt,  will  there- 
fore decompose  spontaneously  in  an  open  vessel.  But  those  with  a 
lower  vapor  tension,  such  as  the  pentahydrate  of  cupric  sulphate  with 
a  tension  of  2  mm.  at  9°,  will  not  do  so.  Granular  calcium  chloride 
CaCl2,2H20  is  used  for  drying  gases,  because  it  has  an  exceedingly 
low  tension  of  water  vapor,  and  combines  with  water  vapor  to  form 
CaCl2,6H2O. 

The  behavior  of  hydrates  does  not  indicate,  as  might  seem  at  first 
sight  to  be  the  case,  that  the  water  is  contained  in  them  in  some  way 
in  the  free  state.  The  fact  is  that  the  above  statements,  with  corre- 
sponding changes  in  the  wording,  might  be  made  of  all  dissociations  in 
chemistry.  Oxides  which  yield  oxygen  when  heated  give  a  different 
pressure  of  oxygen  at  each  temperature,  carbonates  of  carbon  di- 
oxide, and  so  forth. 

The  measurement  of  the  vapor  tension  of  hydrates  gives  definite  information 
in  regard  to  whether  there  are  other  hydrates,  say  of  cupric  sulphate,  with  less 
than  the  normal  number  of  formula-weights  of  water.  If  there  were  only  two 
substances,  CuSO4  and  CuSO4,5H2O,  with  no  compound  of  intermediate  composi- 
tion, then  a  partially  decomposed  specimen  would  be  made  up  partly  of  the  one 
substance  and  partly  of  the  other.  But  if  there  were  an  intermediate  compound, 
say  CuSO4,3H2O,  then  desiccating  a  specimen  of  the  pentahydrate  would  give 
nothing  but  mixtures  of  CuSO4,3H2O  and  CuSO4,5H2O  until  all  the  latter  was 
decomposed.  Then,  and  only  then,  the  trihydrate  would  begin  to  lose  water. 
Now  the  trihydrate,  being  a  definite  and  different  substance,  would  have  a  vapor 
tension  of  its  own,  and  experimental  study  would  show  its  presence. 

Experiment  shows  that  there  really  are  several  hydrated  cupric  sulphates. 
The  pentahydrate,  at  50°,  has  a  vapor  tension  of  47  mm.,  and  this  vapor  tension  is 
observed  so  long  as  any  pentahydrate  remains  to  be  decomposed.  As  soon  as  the 


WATER 


153 


proportion  of  water  goes  down  to  CuSO4,3H2O,  the  vapor  tension  suddenly  drops 
to  30  mm.  As  the  desiccation  continues,  this  tension  is  maintained  until  the  com- 
position has  reached  CuSO4,H2O.  At  this  point  the  vapor  pressure  falls  to  that 
of  the  monohydrate,  4.5  mm.,  and  remains  at  this  value  until  all  the  rest  of  the 
water  has  been  removed.  Had  there  been  no  intermediate  compound  with  3H2O 
the  tension  would  have  dropped  at  once  from  47  mm.  to  4.5  mm.  If,  conversely, 
we  try  to  combine  water  as  vapor  with  anhydrous  cupric  sulphate,  at  50°,  a  vapor 
pressure  of  at  least  4.5  mm.  is  required  to  cause  union  to  take  place.  The  union 
stops  when  one  formula-weight  of  water  has  undergone  combination.  To  intro- 
duce more,  the  concentration  of  the  water  vapor  must  be  increased  to  nearly  seven 
times  its  first  value,  namely,  to  30  mm.  pressure.  This  enforces  combination  up 
to  CuSO4,3H2O.  For  further  hydration,  a  still  higher  pressure  of  water  vapor  is 
needed  (47  mm.),  and  the  absorption  ceases  when  CuSO4,5H2O  has  been  formed. 
There  are  thus  three  distinct  reversible  actions  which  succeed  one  another  as 
the  hydration  proceeds: 

CuSO4  +  H2O  <=±  CuSO4,H2O. 
CuSO4,H2O  +  2H2O  <=±  CuSO4,3H2O. 
CuSO4,3H2O  +  2H2O  +±  CuSO4,5H2O. 

The  first  represents  a  greater  affinity  than  the  second,  and  the  second  than  the 
third. 

The  graphic  representation  of  these  facts  (Fig.  56)  will  make  the  behavior  of 
the  compounds  clearer.  The  proportion  of  water  combined  with  one  formula- 
weight  of  cupric  sulphate  is  laid  off  along  the  horizontal  axis.  The  pressures  at 
which  it  enters  or  leaves  the 
compounds  at  50°  are  the  or- 
dinates.  As  far  as  1H2O  the 
pressure  is  constant  (4.5mm.). 
Beyond  that  point  and  up  to 
3H2O  it  is  constant  but  much 
higher.  Between  3H20  and 
5H2O  it  is  constant  again  but 
higher  still. 

The  tension  of  free  water 
at  the  same  temperature  is  92 
mm.  It  is  constant  irrespective 
of  the  amount  of  water,  and 


60 
47— 


40 


20 


10 


4.5 
0 


1234 

Formula-weights  of  water 

FIG.  56. 


would  therefore  be  on  a  single 
continuous  line  parallel  to  the 
horizontal  axis  and  twice  as 
high  above  it  as  the  uppermost 

one  in  the  diagram.  If,  at  50°,  a  vessel  of  water  were  put  under  a  bell  jar  along- 
side of  anhydrous  cupric  sulphate,  its  vapor  would  be  more  than  sufficiently  con- 
centrated fully  to  hydrate  the  compound.  Again,  while  4.5  mm.  pressure  of 
water  vapor  will  cause  water  to  combine  with  anhydrous  cupric  sulphate  at  50°, 


154  INORGANIC  CHEMISTRY 

a  pressure  of  92  mm.  will  be  required  to  liquefy  the  water  vapor  at  the  same 
temperature. 

The  above  discussion  shows  that  the  last  formula-weight  of  water,  in  hydrates 
which  dissociate  by  stages,  is  not  different  in  kind  from  the  others.  It  differs  only 
in  the  degree  of  tenacity  with  which  it  is  held.  It  is  therefore  unnecessary, 
merely  on  this  account,  to  dignify  it  by  the  separate,  and  misleading  name  of 
water  of  constitution,  as  has  been  done  by  some  chemists. 

Water  of  hydration  is  frequently  called  water  of  crystallization,  on  account 
of  the  fact  that  when  water  is  driven  off  by  heating,  the  substance  usually  crumbles 
to  powder  ( effloresces ).  The  term  is  decidedly  misleading,  however.  It  suggests 
that  water  and  crystallization  are  related  in  some  way,  which  is  not  the  case. 
Sulphur,  galena,  potassium  chlorate,  and  thousands  of  other  crystallized  sub- 
stances, do  not  contain  the  elements  of  water.  Nor  do  the  substances  which 
combine  with  water  remain  amorphous  (without  crystalline  form)  in  its  absence. 
They  all  crystallize  from  the  molten  condition  or  from  some  non-aqueous  solvent, 
although,  as  substances  different  from  the  hydrates,  their  crystalline  form  is 
different.  Iceland  spar,  or  any  other  crystallized  carbonate  which  can  be  de- 
composed by  heating,  becomes  opaque  and  porous  or  falls  to  powder  when  the 
carbon  dioxide  is  driven  out.  But  it  has  not  occurred  to  any  one  to  call  this 
carbon  dioxide  of  crystallization!  The  fact  is  that  all  pure  chemical  substances, 
in  solid  form,  when  in  a  stable  physical  condition,  are  crystalline.  Amorphous 
substances,  like  wax  and  glass,  are  always  supercooled  liquids. 

The  term  arose  from  a  misconception,  and,  when  used,  always  succeeds  in 
transmitting  the  misconception  along  with  the  name.  The  ease  with  which  some 
of  the  hydrates  decomposed  suggested  the  idea  that  they  contained  water  as  a 
discrete  substance.  There  is  no  more  justification  for  this  idea,  however,  than  for 
the  notion  that  carbonates  contain  ready-made  carbon  dioxide.  The  hydrates 
contain  the  elements  of  water  just  as  sugar  and  alcohol  do,  and  there  is  no  evi- 
dence that  they  "contain  water"  in  any  other  sense  than  that  in  which  the  phrase 
might  be  used  of  these  organic  bodies. 

In  consequence  of  their  decomposition  into  and  formation  from  substances 
capable  of  separate  existence,  the  hydrates  are  classed  with  molecular  compounds 
(q.v.).  The  behavior  of  the  compounds  of  salts  with  ammonia  (like  2AgCl,3NH3), 
with  nitric  oxide,  and  with  each  other  (double  salts),  is  quite  similar. 

How  Formulse  and  Equations  are  Obtained.  —  In  the  last 
few  pages  several  formulae  (e.g.,  of  hydrates)  and  several  new  equa- 
tions have  been  given.  How  do  we  know  what  to  set  down  in  mak- 
ing an  equation?  We  cannot  learn  this  by  simply  writing  formulae 
on  a  piece  of  paper.  In  each  case,  experiments  must  be  made  in 
the  laboratory.  For  example,  how  do  we  know  that  the  common 
hydrate  of  cupric  sulphate  has  the  formula  CuSO4,5H2O,  and  not 
CuSO4,H2O?  We  must  make  a  quantitative  experiment.  We  weigh 
a  porcelain  dish  or  crucible,  first  empty,  and  then  with  a  little  of  the 


WATER  155 

hydrate.  Suppose  the  difference  in  weight  to  be  2.05  g.  (=  weight; 
of  hydrate).  We  then  heat  the  dish  and  contents,  until  the  water  is 
driven  out,  and  weigh  again.  The  difference  is  now  only  1.31  g.  (wt 
of  anhydrous  cupric  sulphate).  The  water,  therefore,  weighed  2.05  — 
1.31  =  0.74  g.  Assuming  that  we  know  the  formulae  (compositions) 
of  cupric  sulphate  and  of  water,  we  obtain  their  formula-weights: 
CuSO4  =  63.57  +  32.06  +  4X16  =  159.63;  and  H2O  =  2X1.008  + 
16  =  18.016.  The  formula  must  be  CuSO4,zH20.  Also, 

159.63  :  x  X  18.016  :  :  1.31  :  0.74. 

Solving  for  x  we  have  x  X  18.016  X  1.31  =  159.63  X  0.74,  or  x  = 
159.63  X  0.74/18.016  X  1.31  =  5.00.  The  formula  is,  therefore, 
CuSO4,5H2O,  and  the  equation  for  the  decomposition:  CuSO4,5H20 


To  make  an  equation,  we  must  note  what  substances  are  taken,' 
and  recognize  by  their  properties  all  the  substances  produced.  If  all 
the  substances  are  well  known,  and  we  can  find  their  formulae  in  a 
book,  we  can  make  the  equation  at  once.  If  we  cannot  find  the 
formulae,  we  make  measurements  to  determine  the  proportions  by 
weight,  calculate  the  formulae,  and  then  make  the  equation. 

Composition  of  Water.  —  The  measurement  of  the  propor- 
tions by  weight  and  volume  in  which  hydrogen  and  oxygen  combine 
to  form  water  has  been  the  subject  of  a  larger  number  of  elaborate 
investigations  than  any  other  single  problem  of  this  kind.  The 
difficulty  in  making  the  former  measurement  arises  from  the  fact  that 
both  constituents  are  gases,  and  are  therefore  difficult  to  weigh. 

From  1842  until  1888,  the  accepted  value  of  this  proportion  was  that  of  the 
French  chemist,  Dumas.  His  experiments  gave  the  ratio  of  hydrogen  to  oxygen 
2  :  15.96,  or  2  X  1.0025  :  16.  Cooke  and  Richards  (1888)  obtained  a  figure  for 
the  oxygen  appreciably  smaller  than  this,  namely  2  :  15.90,  or  2  X  1.0063  :  16, 
and  Lord  Rayleigh  (1889)  2  :  15.926.  W.  A.  Noyes  (1889-1907)  made  many 
determinations  and,  finally,  from  twenty-four  measurements  obtained  the  ratio 
2  :  15.875,  or  2  X  1.00787  :  16. 

The  investigation  (lasting  twelve  years)  which  finally  settled  this  question 
was  that  of  Edward  Morley.  The  most  striking  of  his  experiments  consisted 
in  a  series  of  syntheses  of  water,  in  which  he  weighed  the  hydrogen  as  well  as 
the  oxygen,  and  afterwards  weighed  the  water  produced  from  them.  The 
hydrogen  was  confined  by  absorption  in  palladium  (p.  124),  and  could  thus  be 
contained  in  large  quantity  in  a  small,  elongated  bulb.  During  the  progress  of  the 
experiment  it  was  driven  out  by  a  suitable  heating  arrangement.  The  oxygen 
was  contained  in  large  globes  holding  15-20  liters.  The  losses  in  weight  of  the 


156 


INORGANIC  CHEMISTRY 


palladium  tube  and  of  the  globes  gave  the  hydrogen  and  oxygen  consumed.  The 
manipulator  in  which  the  gases  were  combined  and  the  water  collected  is  repre- 
sented in  Fig.  57.  The  gases  entered  through  two  small  tubes  marked  A.  Just 
above  them,  between  two  platinum  wires,  a  discharge  of  electricity  started  the 
union  and  when  necessary  maintained  it.  The  vessel  was  first  filled  by  admitting 
oxygen,  and  the  hydrogen  was  burned  at  the  mouth  of  the  tube  from  which  it 
issued.  This  part  of  the  apparatus  was  immersed  in  a  vessel 
of  water  with  transparent  walls  through  which  the  union  could 
be  watched,  and  the  steam  formed  was  condensed  and  collected 
in  the  bottom  of  the  vessel.  The  vacuum  thus  produced  en- 
abled the  oxygen  continually  to  flow  into  the  manipulator  from 
the  globes.  In  this  way  forty-two  liters  of  hydrogen  and 
twenty-one  liters  of  oxygen  could  be  combined  in  about  an 
hour  and  a  half. 

At  the  end  of  the  experiment  this  part  of  the  apparatus 
was  disconnected  and  placed  in  a  freezing  mixture  which  con- 
verted the  water  into  ice  and  practically  condensed  the  whole 
of  its  vapor.  The  uncombined  gas  in  the  apparatus  was  with- 
drawn and  its  nature  and  quantity  determined.  The  increase 
in  weight  of  the  manipulator  gave  the  quantity  of  water  formed. 
The  success  of  each  experiment  could  be  tested  by  comparing 
the  sum  of  the  weights  of  oxygen  and  hydrogen  with  that  of 
the  water  obtained  from  them.  The  manipulation  was  so 
skilful,  and  the  various  corrections  used  were  so  adequate, 
that  this  difference  was  almost  negligible.  The  ratio  of  hy- 
drogen to  oxygen  in  water  in  this  series  of  experiments  was 
2  :  15.879,  a  result  which  agreed  with  other  methods  of  de- 
termining the  same  ratio  which  Morley  used.  It  agrees  also 
with  the  average  of  the  numbers  obtained  by  other  observers. 


.A 


FIG.  57. 


The  most  probable  value  of  the  ratio  by  weight, 
taking  his  own  and  other  trustworthy  measurements 
into  account,  is  given  by  Morley  as  2  :  15.879,  or 
2.015  :  16.  The  proportion  by  weight  at  present  accepted  is  2  X 
1.008  :  16.  The  proportion  by  volume  is  2.0027  volumes  of  hydrogen 
to  1  volume  of  oxygen. 

That  the  proportion  by  volume  is  very  close  to  2  :  1  may  easily  be 
shown.  We  may  use  a  U-shaped  tube  closed  at  one  end  by  a  stopcock 
and  graduated  (Fig.  58).  At  first,  the  left  limb  of  the  tube,  called  a 
eudiometer,  is  filled  with  mercury.  One  of  the  gases  is  admitted  so  as 
to  fill  a  portion  of  the  tube  and,  the  levels  having  been  equalized  (cf. 
p.  105),  the  volume  of  the  gas  is  read.  Then  some  of  the  other  gas  is 
introduced  and  the  leveling  and  reading  repeated.  Let  us  suppose 
that  15  c.c.  of  hydrogen  and  10  c.c.  of  oxygen  have  thus  been  taken. 


WATER 


157 


The  right  limb  is  then  filled  with  mercury  and  closed  firmly  with  the 
thumb.  A  spark  from  an  induction  coil  passing  between  the  two 
short  platinum  wires  near  the  top  of  the  tube  explodes  the  mixture. 
The  steam  produced  by  the  union  condenses  almost  immediately 
and  occupies  practically  no  volume  worth  considering.  When  the 
thumb  is  removed,  the  mercury  rises  on  the  left  and  fills  up  the  space 
left  by  the  disappearance  of  part  of  the  gases.  Unless 
the  proportion  taken  happens  to  have  been  exact, 
some  of  one  or  other  of  the  gases  will  remain.  Its 
volume  is  measured  by  equalizing  the  levels  and  read- 
ing as  before.  In  the  case  we  have 
imagined,  the  residual  gas  is  oxy- 
gen, and  there  are  almost  exactly 
2.5  c.c.  of  it.  It  is  evident,  there- 
fore, that  15  c.c.  of  hydrogen 
united  with  7.5  c.c.  of  oxygen;  in 
other  words,  the  proportion  by 
volume  is  2  :  1. 

Gay-Lussac9s  Law  of  Com- 
bining Volumes.  —  The  almost 
mathematical  exactness  with 
which  small  integers  express  this 
proportion  is  not  a  mere  coinci- 
dence. Whenever  gases  unite,  or 
gaseous  products  are  formed,  the 
proportions  by  volume  (measured 
at  the  same  temperature  and  pres-  ' 
sure)  of  all  the  gaseous  bodies  con- 
cerned can  be  represented  very 
accurately  by  ratios  of  small  in- 
tegers. This  is  called  Gay-Lussac's  law  of  combining  volumes  (1808). 
Thus,  when  the  above  experiment  is  carried  out  at  100°,  in  order  that 
the  product,  water,  may  be  gaseous  also,  it  is  found  that  the  three 
volumes  of  the  constituents  give  almost  exactly  two  volumes  of  steam. 
For  example,  15  c.c.  of  hydrogen  and  7.5  c.c.  of  oxygen  give  15  c.c. 
of  steam.  Of  course  the  hydrogen,  oxygen,  and  steam  must  be 
measured  at  the  same  pressure,  and  the  temperature  must  remain 
constant  (100°)  during  the  experiment.  Proper  manipulation  secures 
the  former,  and  a  jacket  filled  with  steam  (Fig.  59)  the  latter  condi- 
tion. Strips  of  paper,  1,  2,  and  3,  are  pasted  on  the  jacket  in  such  a 


FIG.  58. 


FIG.  59. 


158  INORGANIC  CHEMISTRY 

way  that  equal  lengths  of  the  eudiometer,  in  this  case  a  straight  one, 
are  laid  off.  The  three  divisions  having  been  filled  with  a  mixture  of 
hydrogen  and  oxygen  in  the  proper  proportions,  the  gas,  after  the  ex- 
plosion, shrinks  so  as  to  occupy,  at  the  same  pressure,  only  two  of  them. 
From  this  universal  truth  in  regard  to  the  combination  of  gases, 
we  draw  the  important  inference  that  the  chemical  unit-weights  of 
simple  substances,  and  the  formula-weights  of  compounds,  in  the 
gaseous  condition,  occupy  at  the  same  temperature  and  pressure  vol- 
umes which  are  equal  or  stand  to  one  another  in  the  ratio  of  small 
integers  (see  Molar  weights). 

The  chemical  behavior  of  the  other  compound  of  hydrogen  and  oxygen, 
hydrogen  peroxide  (q.v.),  is  difficult  to  comprehend  until  further  experience  has 
been  gained.  Then,  too,  its  formula  H2O2  cannot  be  justified  until  the  means  of 
determining  molar  weights  in  solution  have  been  discussed.  In  view  of  these 
facts,  it  will  be  taken  up  later. 

Exercises.  —  1.  Name  some  other  transitions  from  one  physical 
state  to  another  which  are  familiar  (p.  144). 

2.  What  evidence  is  there  in  the  common  behavior  of  ether  and 
chloroform  to  show  that  these  liquids  have  high  vapor  tensions? 

3.  If  the  pressure  of  the  steam  in  a  boiler  is  ten  atmospheres,  at 
what  temperature  is  the  water  boiling  (p.  146)? 

4.  How  many  grams  of  water  could  be  heated  from  20°  to  100°  by 
the  heat  required  to  melt  1  kgm.  of  ice  at  0°? 

5.  What  do  you  infer  from  the  fact  that  alum  and  washing  soda 
lose  their  water  of  crystallization  when  left  in  open  vessels,  while 
gypsum  does  not  (p.  152)? 

6.  Which  facts  show  most  conclusively  that  hydrates  are  true 
chemical  compounds? 

7.  In  what  ways  does  a  hydrate  differ  from,  (a)  a  solution,  (6)  an 
hydroxide? 

8.  Should  you  expect  to  find  any  difference,  in  respect  to  chemical 
activity,  between  the  three  forms  of  water  (ice,  water,  and  steam)? 
If  so,  arrange  them  in  the  order  of  probable  increasing  activity  (pp. 
93-94).     Have  we  had  any  experimental  confirmation,  or  the  reverse, 
of  this  conclusion? 

9.  Which  contains  more  internal  energy,  and  is  therefore  more 
active,  the  anhydrous  substance,  or  the  corresponding  hydrate? 

10.  Gypsum  is  a  hydrate  of  calcium  sulphate  CaSO4.     If  6  g.  of 
gypsum,  when  heated,  lose  1.256  g.  of  water,  what  is  the  formula  of 
the  hydrate? 


WATER  159 

11.  At  what  temperature  will  ice  melt  in  a  vacuum? 

12.  Indicate  briefly  the  objections  to  the  form  of  statement 
(commonly  used) :  "Blue  vitriol  loses  4  molecules  of  water  of  crystal- 
lization at  100°  and  the  fifth  at  200°."     Re-write  this  correctly. 

13.  Equal  weights  (say,  249.7  g.)  of  blue-stone  and  of  CuS04  are 
inclosed  at  50°,  occupying  the  whole  volume.     Find  the  composition 
of  the  system  at  equilibrium. 


CHAPTER  IX 

RELATIONS   BETWEEN   THE   STRUCTURE  AND   BEHAVIOR  OP 
MATTER.     THE  KINETIC -MOLECULAR  VIEWPOINT 

WE  have  seen  (p.  65)  that  matter  is  composed  of  minute  particles 
called  molecules.  Just  as  we  can  thoroughly  understand  the  be- 
havior of  a  watch  or  an  automobile  engine  only  if  we  know  the  details 
of  its  structure,  and  how  the  parts  work,  so  we  can  understand  the 
physical  and  chemical  behavior  of  matter  in  masses  only  if  we  are 
familiar  with  its  ultimate  mechanism.  Hence,  we  must  now  take  up 
the  structure  of  matter  in  its  three  states,  the  gaseous,  the  liquid,  and 
the  solid.  In  doing  this,  we  shall  keep  constantly  in  view  the  con- 
nection between  the  molecular  relations  and  the  general  behavior  of 
the  matter. 

The  Molecular  Structure  of  Gases.  —  The  most  noticeable 
fact  about  gases  is  that  they  can  be  compressed  to  an  enormous  extent. 
Oxygen  at  760  mm.,  for  example,  can  be  reduced  by  pressure  to  one 
two-hundredth  of  its  volume,  or  even  less.  The  compression  does 
not  affect  the  individual  molecules,  and  therefore  does  not  diminish 
the  volume  actually  occupied  by  the  oxygen,  but  crowds  the  mole- 
cules closer  together  and  diminishes  to  one  two-hundredth  the  space 
between  them.  Compressing  a  gas  is,  in  fact,  mainly  reducing 
the  empty  space  of  which  it  chiefly  consists.  To  understand  what 
follows,  the  reader  must  keep  constantly  and  vividly  before  him  a 
mental  image  of  a  jar  of  gas  as  consisting  of  small  particles  separated 
by  relatively  wide,  empty  spaces.  The  molecules  are  in  rapid  motion 
and  move  in  straight  lines,  excepting  when  they  strike  one  another 
or  the  wall  of  the  vessel,  and  rebound. 

The  actual  number  of  molecules  in  a  given  volume  of  a  gas  will 
be  discussed  in  Chapter  XII,  where  the  figures  can  be  given  in  a 
more  memorable  form. 

The  Qualitative  Properties  of  Gases.  —  Let  us  now  note  the 
more  obvious  qualities  of  gases,  printing  in  italics  the  fact  concerning 
a  mass  of  gas,  and  in  black  type  the  property  of  the  molecules  which 
accounts  for  the  fact. 

160 


THE  KINETIC-MOLECULAR  VIEWPOINT  161 

The  most  remarkable  thing  about  a  gas,  considering  the  looseness 
with  which  its  material  is  packed,  is  the  total  absence  in  it  of  any 
tendency  to  settling  or  subsidence.  Since  the  molecules  cannot  be  at 
rest  upon  one  another,  as  the  great  compressibility  shows,  we  are  driven 
to  suppose  that  they  are  widely  separated  from  one  another,  and  that 
they  occupy  the  space  by  constantly  moving  about  in  all  directions. 
But  a  moving  aggregate  of  particles  which  does  not  even  finally  settle 
must  be  in  perpetual  motion.  We  must,  therefore,  imagine  the  mole- 
cules to  be  wholly  unlike  visible  particles  of  matter  in  having  perfect 
elasticity,  in  consequence  of  which  they  undergo  no  loss  of  energy  after 
a  collision.  They  must  continually  strike  the  walls  of  the  vessel  and 
one  another  and  rebound,  yet  without  loss  of  motion.  The  fact  that 
each  gas  is  homogeneous,  efforts  to  sift  out  lighter  or  heavier  samples 
having  failed,  requires  the  supposition  that  all  the  molecules  of  a  pure 
gas  are  closely  alike. 

The  diffusibility  of  gases  is  due  to  the  motion  of  the  molecules, 
and  their  permeability  to  the  space  available  to  receive  molecules  of 
another  gas.  These  two  modes  of  behavior  involve  no  additional 
molecular  properties.  The  word  "diffusion"  is  often  thought  to 
mean  the  property  of  a  given  mass  of  gas  in  virtue  of  which  another 
gas  can  mix  with  the  given  mass.  This  property  is  not  diffusibility 
but  permeability.  It  is  the  motion  of  each  gas,  making  its  way  into 
the  other  gas,  which  is  diffusion.  Diffusion  is  spontaneous  motion  of 
the  parts  of  a  gas  away  from  their  original  location.  Unless  this 
motion  is  into  an  empty  space,  the  diffusing  molecules  must,  of  course, 
move  into  another  body  of  gas.  In  the  case  of  the  jars  of  carbon 
dioxide  and  air  (p.  125),  each  gas  moved  in  part  out  of  its 
original  jar  (diffused),  and  each  received  parts  of  the 
other  gas  into  its  jar  (was  permeated). 


I 


Quantitative  Properties  of  Gases:  Boyle9 s  and 
Charles9  Laws.  —  Passing  now  to  Boyle's  law  (p.  106), 
the  thing  to  be  accounted  for  is  that  when  a  sample  of 
a  gas  diminishes  in  volume,  its  pressure  increases  in  the 
same  proportion.  Let  the  diagram  (Fig.  60)  represent  a 
cylinder  with  a  movable  piston,  upon  which  weights  may 
be  placed  to  resist  the  pressure.  Now  the  pressure  ex- 
ercised by  the  gas  cannot  be  explained  as  being  like  the  pressure  of 
the  hand  upon  a  table,  since  we  have  just  assumed  that  the  particles 
are  not  even  approximately  at  rest,  and  the  spaces  between  them  are 
enormous  compared  with  the  size  of  the  molecules  themselves.  The 


FIG.  60. 


162  INORGANIC  CHEMISTRY 

gaseous  pressure  must  therefore  be  attributed  to  the  colossal  hailstorm 
which  their  innumerable  impacts  upon  the  piston  produce.  If  this  is 
the  case,  the  compressing  of  a  gas  must  consist  simply  in  moving 
the  partition  downwards  so  that  the  particles  as  they  fly  about  are 
gradually  restricted  to  a  smaller  and  smaller  space.  Their  paths 
become  on  an  average  shorter  and  shorter.  Their  impacts  upon  the 
wall  become  more  and  more  frequent.  So  the  pressure  which  this 
occasions  becomes  greater  and  greater,  and  is  proportional  to  the 
degree  of  crowding  (the  concentration)  of  the  molecules. 

There  are  two  other  points  which  must  be  added.  When  we 
diminish  the  volume  to  one-half,  we  find  from  experience  that  the 
pressure  becomes  exactly,  or  almost  exactly,  twice  as  great.  This 
must  mean  that  although  the  particles  are  becoming  crowded  they 
do  not  interfere  with  one  another's  motion,  excepting  of  course  where 
actual  collision  causes  a  rebound.  Only  in  the  absence  of  interference 
would  doubling  the  number  of  particles  per  unit  of  volume  give 
exactly  double  the  number  of  impacts  on  the  walls.  Hence  the 
particles  must  have  practically  no  tendency  to  cohesion.  Again,  the 
molecules  must  move  in  straight  lines,  because,  if  they  moved  in 
orbits  of  some  kind,  many  of  the  orbits  would  not  be  intersected  by 
the  wall  of  the  vessel  until  great  reduction  in  the  volume  had  taken 
place,  and  thus,  as  the  volume  diminished,  the  frequency  of  the 
impacts,  and  therefore  the  pressure,  would  increase  faster  than  the 
concentration. 

Boyle's  law  therefore  adds  four  more  conceptions  to  our  molecular 
hypothesis,  namely,  that  the  impacts  of  the  molecules  produce  the 
pressure,  that  the  crowding  of  the  molecules  represents  the  concentra- 
tion (p.  103),  and  that  the  particles  move  in  straight  lines  and  show 
almost  no  cohesion,  since  pressure  and  concentration  are  very  closely 
proportional  to  one  another. 

It  will  be  seen,  on  consideration,  that  if  the  molecules  are  assumed  to  repel 
one  another,  they  would  do  so  more  violently  the  more  closely  they  were  packed 
together.  This  assumption  would  therefore  suit  the  case  of  a  gaseous  body  in 
which  the  pressures  increased  according  to  some  power  of  the  concentration  other 
than  the  first,  and  therefore  much  more  rapidly  than  in  known  gases.  In  spite  of 
its  inapplicability,  this  notion  is  supposed  by  many  people  to  be  part  of  the 
kinetic  theory. 

Charles'  law  (p.  110),  that  a  gas  receives  equal  increments  in  volume 
or  pressure  for  equal  elevations  in  temperature,  requires  but  one  addi- 
tion to  the  hypothesis.  Concretely,  if  our  specimen  of  gas  (Fig.  60)  is 


THE  KINETIC-MOLECULAR  VIEWPOINT  163 

at  0°,  and  we  permit  its  pressure  to  remain  constant  by  leaving  the 
same  weight  on  the  piston,  then  when  the  temperature  of  the  gas  is 
raised  to  1°,  the  volume  will  gain  ^73  of  the  original  volume.  If,  on 
the  other  hand,  we  restrict  the  gas  to  the  original  volume,  the  pressure 
will  evidently  increase,  and  the  augmentation  will  be  *H  of  the 
original  pressure.  Now,  how  can  we  account  for  an  increase  in 
pressure  as  the  result  of  heating  a  mass  of  rapidly  moving  molecules? 
The  action  of  a  particle  colliding  with  a  surface  is  measured  in  physics 
in  terms  of  its  mass  and  its  velocity.  It  is  evident  that  heating  a  cloud 
of  molecules  would  not  increase  the  mass  of  each,  and  it  must  there- 
fore increase  the  velocity  of  each,  since  the  kinetic  energy  of  all 
becomes  greater.  This  conclusion  is  in  harmony  with  our  experience 
that  violently  rubbing  a  solid  raises  its  temperature,  and  such  a  mode 
of  treatment  might  plausibly  be  supposed  to  communicate  motion  to 
the  minute  parts  of  the  body. 

Gay-Lussac's  and  Avogadro's  Laws.  —  The  fact  that  the  com- 
bining volumes  of  gaseous  substances  are  equal,  or  stand  to  one  another 
in  the  ratio  of  small  whole  numbers  (cf.  p.  157),  suggests  two  ideas: 
First,  that  chemical  combination,  considered  in  detail,  and  arranged 
to  harmonize  with  this  theory,  would  involve  unions  of  a  few 
particles  of  more  than  one  kind  to  form  composite  molecules.* 
And,  second,  that  a  simple  integral  relation  must  be  assumed  to  exist 
between  the  numbers  of  molecules  in  equal  volumes  of  different  gases, 
•  at  the  same  temperature  and  pressure.  Avogadro  (1811),  the  pro- 
fessor of  physics  in  Turin,  put  forward  the  hypothesis  that  these 
numbers  might  be  equal.  A  more  strict  study  of  the  assumptions  we 
have  been  making,  and  of  some  additional  facts,  has  since  shown  that 
no  other  conjecture  than  Avogadro's  would  be  consistent  with  them. 
Thus  it  now  bears  the  relation  of  a  logical  deduction  from  the  kinetic- 
molecular  theory  and  the  properties  of  gases,  and  is  known  as  Avo- 
gadro's law.  It  may  also  be  put  in  the  form:  At  the  same  tempera- 
ture and  pressure,  the  molecular  concentration  (cf.  p.  104)  of  all  kinds 
of  gases  has  the  same  value. 

Diffusion.  —  The  law  of  diffusion  (p.  126)  harmonizes  with  the 
kinetic-molecular  theory  without  further  modification  of  the  latter. 
The  strict  deduction  of  this  law,  as  well  as  of  the  preceding  ones,  from 

*  This  is  essentially  the  idea  used  by  Dalton  (p.  65),  before  Gay-Lussac's 
law  was  known,  however,  for  the  explanation  of  the  laws  of  chemical  combina- 
tion. 


164  INORGANIC  CHEMISTRY 

our  series  of  assumptions,  will  be  found  in  any  work  on  physical 
chemistry. 

The  speed  of  the  hydrogen  molecule  at  room  temperature  is  1840 
meters  per  second.  The  masses  of  the  hydrogen  and  oxygen  mole- 
cules are  as  1  :  16,  and  the  speeds  of  diffusion  (p.  126)  as  Vl6  :  VI, 
or  4  :  1.  Hence  the  speed  of  the  oxygen  molecule  is  one-fourth  of 
1840,  or  460  m.  per  sec. 

Calculation  shows  the  activity  of  the  molecules  to  be  so  great 
that,  in  air,  the  number  striking  a  single  square  centimeter  of 
surface  every  second  would  fill  no  less  than  twenty  liters. 

Two  Deviations  front  Boyle  s  Law.  —  Finally,  we  have  re- 
ferred (p.  108)  to  the  fact  that  at  low  pressures  the  concentration  increases 
more,  and  at  high  pressures  much  less  than  Boyle's  law  indicates.  The 
former  effect  is  brought  into  accord  with  our  hypothesis  when  we 
remember  that  the  matter  even  of  gases  can  cohere,  as  is  shown 
plainly  when  gases  are  solidified.  The  tendency  of  the  molecules  to 
cohere  must  therefore  show  itself  in  the  gaseous  condition  by  pulling 
the  gas  together  and  producing  somewhat  greater  concentration  than 
is  strictly  consistent  with  the  value  of  the  pressure.  Thus,  2  liters  of 
oxygen  at  760  mm.  and  0°,  when  subjected  to  2  atmospheres  pressure, 
give  0.9991  liters  instead  of  1  liter.  The  additional  contraction  of 
0.0009  liters  (0.9  c.c.)  is  due  to  the  effect  of  the  cohesion  when  the 
molecules  are  thus  crowded  closer  together.  The  gases  which  are 
more  easily  liquefied  than  is  oxygen  show  greater  effects.  Thus,  2 
liters  of  sulphur  dioxide  at  760  mm.  and  0°,  when  subjected  to  2 
atmospheres  pressure,  give  only  0.974  liters,  showing  a  contraction 
due  to  cohesion  of  26  c.c.  At  temperatures  below  0°  the  contractions 
due  to  cohesion  become  rapidly  greater.  This  cohesion  is  not  of  the 
nature  of  gravitational  attraction.  Although  this  effect  of  cohesion 
is  usually  insignificant,  the  modern  method  of  liquefying  gases  (q.v.) 
depends  upon  it  almost  entirely. 

The  abnormally  small  reductions  in  volume  which  occur  when  the 
volume  of  the  gas  has  already  been  greatly  reduced  remind  us  that, 
according  to  our  hypothesis,  it  is  only  the  space  between  the  molecules 
that  is  diminished  as  pressure  rises,  and  not  the  space  occupied  by  the 
molecules.  Hence,  when  the  molecules  have  become  so  crowded  to- 
gether that  this  irreducible  space  begins  to  form  an  appreciable 
fraction  of  the  whole,  a  doubling  of  the  pressure  will  diminish  to 
one-half  its  value  only  a  part  (the  vacant  part)  of  the  volume  the  gas 
occupies. 


van 


THE  KINETIC-MOLECULAR  VIEWPOINT  165 

If  the  incompressible  space  occupied  by  the  molecules  is  called  6,  and  that  of 
the  whole  gas  v,  then  the  amended  form  of  Boyle's  law  reads  p  (v  —  6)  =  constant. 
Similarly,  if  the  cohesive  tendency  is  taken  into  account,  it  is  plain  that  its  effect 
will  be  numerically  greater  at  small  volumes,  although  not  so  easily  observed.  It 
is  in  fact  inversely  proportional  to  the  square  of  the  volume.  If  it  is  expressed 
in  the  same  units  as  the  pressure  by  a,  the  total  of  the  compressing  tendencies 

becomes  p  -+•  — .     Hence  Boyle's  law,  for  constant  temperatures,  as  amended  by 

\r 

der  Waals,  reads  [p  +  ~  }  (v  —  b)  =  constant,  a  formula  which  describes 

the  actual  behavior  of  most  gases  with  remarkable  accuracy.  Hydrogen  alone,  at 
ordinary  temperatures,  shows  no  excessive  compressibility  at  low  pressures. 
Thus,  2  liters  of  hydrogen  at  760  mm.,  when  subjected  to  1520  mm.  pressure, 
give  1.0006  liters,  or  0.6  c.c.  more  than  Boyle's  law  suggests.  The  cohesion  (a) 
is  here  negative,  so  that  the  effect  of  the  constant  (6)  prevails  from  the  very 
first. 

It  will  be  noted  that  the  kinetic  theory  affords  a  mechanical 
explanation  (i.e.,  description  in  detail)  of  all  the  properties  of  gases, 
with  the  exception  of  the  cohesion.  It  includes  no  machinery  for 
explaining  how  the  attraction  operates. 

In  consequence  of  these  two  deviations,  there  are  not  exactly  equal 
numbers  of  molecules  in  equal  volumes  of  any  two  different  gases,  at 
the  same  temperature  and  pressure.  An  imaginary  gas,  which 
exhibits  neither  deviation,  called  a  perfect  gas,  is  often  referred  to  in 
discussing  the  behavior  of  gases. 

Summary.  —  We  may  now  summarize  the  principal  facts  about 
gases  in  mass,  appearing  in  italics  above,  with  the  corresponding 
features  of  the  molecular  relations,  in  heavy  type,  which  we  have 
added  one  by  one. 


FACTS  ABOUT  GASES  IN  MASS. 


CORRESPONDING  RELATIONS  OF  MOLECULES. 


Compressibility 
Diffusibility 

Permeability  . 

Non-settling  . 

Homogeneity  . 

Pressure     .    .  . 

Boyle's  law  .  . 


Vacuum  -f  molecules  widely  separated. 

Molecules  in  rapid  motion. 

Empty  space  relatively  large. 

Molecules  perfectly  elastic. 

Molecules  of  any  one  substance  closely  alike. 

Due  to  impacts  of  molecules. 

Pressure  proportional  to  concentration  of  the 
molecules.  Molecules  move  in  straight 
lines  and,  when  widely  scattered,  show 
almost  no  tendency  to  cohesion. 


166 


INORGANIC  CHEMISTRY 


FACTS  ABOUT  GASES  IN  MASS. 


CORRESPONDING  RELATIONS  OF  MOLECULBS. 


Dalton's  law  (p.  Ill) 
Henry's  law      ... 
Charles'  law 


Gay-Lussac's  law  and  above 
and  other  facts. 


Law  of  diffusion. 

Abnormal  compressibility,  es- 
pecially at  low  pressures. 

Abnormal  incompressibility, 
especially  at  high  pressures. 


Nothing  new. 
Nothing  new. 

Rise  in  temperature  increases  the  velocity, 
and  therefore  the  kinetic  energy  of  the 
molecules. 

There  are  equal  numbers  of  molecules  in 
equal  volumes  of  different  gases  at  the 
same  temperature  and  pressure  (Avoga- 
dro's  law). 

Nothing  new. 

The  tendency  to  cohesion  gives  measurable, 
though  slight  effects. 

The  molecules  themselves  are  incompres- 
sible. 


Critical  Temperature.  —  When  the  concentration  of  a  gas  at 
ordinary  temperatures  is  greatly  increased  by  compression,  the 
cohesive  forces  have  an  opportunity  to  produce  liquefaction.  In 
many  cases,  as  with  sulphur  dioxide  and  carbon  dioxide,  when  the 
approximation  of  the  molecules  has  reached  a  certain  point,  the  liquid 
begins  to  form  on  the  sides  of  the  vessel.  The  condition  is  then 
exactly  the  same  as  that  of  aqueous  vapor  and  water  (p.  145),  and  no 
further  increase  in  pressure  is  required  to  complete  the  liquefaction  of 
the  whole.  The  only  difference  between  steam,  at  a  pressure  below 
the  aqueous  tension  of  water  at  10°,  and  carbon  dioxide  at  the  same 
temperature,  is  that  not  more  than  9.1  mm.  of  pressure  is  required  to 
liquefy  the  steam,  while  about  50  atmospheres  are  needed  to  liquefy 
the  carbon  dioxide. 

There  are  some  gases  in  which,  at  the  ordinary  temperature,  even 
with  the  closest  approximation  of  the  molecules,  the  cohesion  is  unable 
to  overcome  the  motion  of  the  molecules  and  draw  the  material 
together  into  the  more  compact  liquid  form.  Such  gases  are  hydro- 
gen, oxygen,  nitrogen,  and  air,  which  is  a  mixture  of  the  last  two. 
The  remedy  is  obvious.  We  know  of  no  way  to  increase  the  intrinsic 
cohesiveness  of  the  material,  but  we  can  reduce  the  kinetic  energy  of 
the  molecules  by  lowering  the  temperature  of  the  gas.  When  this  has 
been  done  sufficiently,  compression  is  followed  by  liquefaction.  Now 
it  is  found  that  there  is  a  critical  value  for  each  individual  gas,  to  or 
beyond  which  the  kinetic  energy  must  be  reduced  by  lowering  the 
temperature,  before  the  cohesive  tendency  of  that  particular  gas  can 
become  effective  to  produce  liquefaction.  In  1869  Andrews  found 


THE  KINETIC-MOLECULAR  VIEWPOINT  167 

that  carbon  dioxide  could  be  liquefied  at  0°  by  38  atmos.  pressure, 
and  at  30°  by  71  atmos.,  but  that  above  31.35°  it  could  not  be  liquefied 
by  any  pressure.  He  discovered  that  each  gas  has  a  critical  tempera- 
ture, as  he  called  it.  For  carbon  dioxide,  this  temperature  can  be 
observed  by  placing  a  heavy  walled,  glass  tube  (Fig.  61), 
half-filled  with  liquid  carbon  dioxide,  in  a  beaker  of  water, 
and  gradually  raising  the  temperature  of  the  latter.  At 
31.35°,  the  surface  between  the  liquid  and  gas  becomes  hazy 
and  vanishes.  The  vapor  pressure  in  the  gas  has  become 
so  great  that  the  gas  has,  at  this  point,  the  same  density  as 
the  liquid.  When  the  temperature  falls  once  more,  the 
surface  reappears  at  31.35°. 

For  oxygen  this  temperature  is   —118°,  for  hydrogen 
-234°,  for  nitrogen  - 146°.     For  carbon  dioxide  it  is  31.35°, 
for  sulphur  dioxide  156°,  for  water  358°.     The  temperature 
of  a  room  being  below  the  critical  points  of  the  last  three 
substances,  they  are  all  liquefiable  without  cooling,  and  more  easily 
the  farther  the  ordinary   (say  20°),   lies   below  the  critical  tem- 
perature. 

History  of  the  Kinetic- Molecular  Theory.  —  This  theory 
was  first  suggested  by  Daniel  Bernoulli  (1738),  who  explained  by  its 
means  the  pressure  and  compressibility  of  gases.  Lomonossov 
(1748)  developed  the  theory  very  completely  and  explained  Boyle's 
law  and  the  effects  of  changes  in  temperature  by  means  of  it.  He 
also  anticipated  from  the  theory  the  existence  of  the  second  deviation 
from  the  law  of  gases  (1749),  a  discovery  usually  credited  to  Dupre 
(1864).  He  likewise  pointed  out  that  there  was  no  limit  to  the 
maximum  velocity  of  a  molecule,  and  therefore  no  upper  limit  of 
temperature,  but  that  there  must  be  a  lower  limit  (the  absolute  zero) 
at  which  the  molecules  would  be  at  rest  (1744).  This  work  was 
entirely  forgotten,  until  attention  was  called  to  it  in  1904  by  Men- 
schutkin. 

Similar  views  were  formulated  in  detail  by  Waterston  (1845),  but 
were  still  so  much  ahead  of  the  time  that  the  committee  of  the  Royal 
Society  did  not  approve  the  paper  for  publication,  and  it  was  dis- 
covered in  the  archives  of  the  society,  long  afterwards,  by  Lord 
Rayleigh.  The  development  of  the  theory,  so  far  as  it  applies  to 
heat  is  therefore  usually  credited  to  Joule  (1855-60)  and,  in  respect 
to  all  properties  of  gases,  to  Kronig  (1856)  and  Clausius  (1857),  who 
knew  nothing  of  the  earlier  work. 


168 


INORGANIC  CHEMISTRY 


Molecular  Relations  in  Liquids.  —  The  fact  that  even  great 
pressures  produce  little  diminution  in  the  volume  of  a  liquid  shows 
that  the  free  space,  present  in  gases,  is  absent  in  liquids.  The 
measured  effects  of  various  pressures  show,  for  example  in  the  case 
of  water,  that  to  reduce  the  volume  to  one-half  would  require,  not 
doubling  the  pressure  as  in  a  gas,  but  increasing  it  from  1  to  10,000 
atmospheres.  The  molecules  of  a  liquid  are  actually  in  contact  with 
one  another,  and  are  themselves  compressible,  though  with  difficulty 
(Richards) . 

The  phenomena  connected  with  surface  tension,  such  as  coherence 
into  drops,  show  that  cohesion  plays  a  larger  part  in  liquids  than  in 


We  are  further  compelled  to  suppose  that  there  is  motion  of  the 
molecules  inside  the  liquid.  When  alcohol,  as  the  lighter  liquid,  is 
floated  upon  water  in  a  cylinder  (Fig.  62),  the  plane 
separating  the  liquids  is  at  first  easily  visible.  But 
soon  it  becomes  obliterated.  The  water  diffuses  up- 
ward, and  the  alcohol  downward,  each  sifting  its  way 
through  the  other  in  spite  of  gravity.  The  complete 
mixing  of  the  liquids  takes  a  very  much  longer  time 
than  in  the  case  of  two  gases.  It  may  take  months. 
But  even  here  the  hypothesis  helps  us,  by  pointing  to 
the  vast  impediment  which  the  close  packing  of,  and 
therefore  enormous  friction  between  the  molecules  must 
place  in  the  way  of  the  progress  of  any  one  molecule. 
Further,  once  the  mixture  is  formed,  no  tendency  to 
spontaneous  separation  is  ever  observed.  Here  again, 
the  hypothesis  shows  that  none  is  to  be  expected.  If  it 
occurred,  it  would  be  immediately  undone  by  diffusion. 


Fia.  62. 


Molecular  Relations  of  Liquid  and  Vapor. — 

When  the  water  was  introduced  above  the  barometric 
column  (Fig.  53,  p.  145),  the  escape  of  vapor,  that  is,  of  part  of  the 
liquid  in  gaseous  form,  could  have  resulted  only  from  the  spontaneous 
motion  of  the  molecules  in  the  liquid.  Some  of  the  molecules,  moving 
near  the  surface,  went  off  into  the  space  above  the  water  and  became 
gaseous.  To  be  consistent,  we  must  also  conclude  that  the  vapor 
above  the  water  is  not  composed  of  the  same  set  of  molecules  one 
minute  as  it  was  during  the  preceding  minute.  Their  motions  must 
cause  many  of  them  to  plunge  into  the  liquid,  while  others  emerge 
and  take  their  places.  When  the  water  is  first  introduced,  there 


THE  KINETIC-MOLECULAR  VIEWPOINT  169 

are  no  molecules  of  vapor  in  the  space  at  all,  so  that  emission 
from  the  water  predominates.  The  pressure  of  the  vapor  increases 
as  the  concentration  of  the  molecules  of  vapor  becomes  greater, 
hence  the  mercury  column  falls  steadily.  At  the  same  time  the 
number  of  gaseous  molecules  plunging  into  the  water  per  second 
must  increase  in  proportion  to  the  degree  to  which  they  are  crowded 
in  the  vapor.  The  rate  at  which  molecules  return  to  the  water 
thus  begins  at  zero,  and  increases  steadily;  the  rate  at  which  mole- 
cules leave  the  water,  however,  maintains  a  constant  value,  be- 
cause the  properties  of  the  water  are  not  affected  by  changes  in 
the  concentration  of  the  vapor.  Hence  the  rate  at  which  vapor 
molecules  enter  the  water  must  eventually  equal  that  at  which  other 
molecules  leave  the  liquid.  At  this  point,  occasion  for  visible  change 
ceases  and  the  mercury  comes  to  rest.  We  are  bound  to  think, 
however,  of  the  exchange  as  still  going  on,  since  nothing  has  occurred 
to  stop  it.  The  condition  is  not  one  of  rest,  but  of  rapid  and  equal 
exchange.  Such,  described  in  terms  of  the  theory,  is  the  state  of 
affairs  which  is  characteristic  of  a  condition  of  equilibrium  (p.  146). 
The  condition  is  kinetic,  and  not  static. 

Equilibrium.  —  This  term  is  used  so  often  in  chemistry,  and 
is  used  in  so  unfamiliar  a  sense,  that  the  reader  should  consider 
attentively  what  it  implies.  Three  things  are  characteristic  of  a 
state  of  equilibrium: 

1.  There  are  always  two  opposing  tendencies  which,  when  equi- 
librium is  reached,  balance  each  other.     In  the  foregoing  instance, 
one  of  these  is  the  hail  of  molecules  leaving  the  liquid,  which  is 
constant  throughout  the  experiment.      It  represents  the  vapor  ten- 
sion of  the  liquid.      The  other  is  the  hail  of  returning  molecules, 
which,  at  first,  increases  steadily  as  the  concentration  of  the  vapor 
becomes  greater.     This  is  the  vapor  pressure  of  the  vapor.      These 
have  the  effect  of  opposing  pressures  and,  when  the  latter  becomes 
equal  to  the  former,   equilibrium  is  established.     In  all  cases  of 
equilibrium  we  shall  symbolize  the  two  opposing  tendencies  by  two 
arrows,  thus: 

Water  (liq.)  ?=±  Water  (vapor). 

2.  Although  their  effects  thus  neutralize  each  other  at  equilib- 
rium, both  tendencies  are  still  in  full  operation.     In  the  case  in  point, 
the  opposing  hails  of  molecules  are  still  at  work,  but  neither  can 
effect  any  visible  change  in  the  system.     Equilibrium  is  a  state, 
not  of  rest,  but  of  balanced  activities. 


170  INORGANIC  CHEMISTRY 

3  (and  this  is  the  chief  mark  of  equilibrium).  A  slight  change  in 
the  conditions  produces,  seldom  a  great  or  sharp  change,  but  always 
and  instantly,  a  corresponding  small  change  in  the  state  of  the 
system.  The  change  in  the  conditions  accomplishes  this  by  favoring 
or  disfavoring  one  of  the  two  opposing  tendencies.  Thus,  for  ex- 
ample, when  the  temperature  of  a  liquid  is  raised,  the  kinetic  energy 
of  its  molecules  is  increased,  the  rate  at  which  they  leave  its  surface 
becomes  greater,  the  vapor  tension  increases  and,  hence,  a  greater 
concentration  of  vapor  can  be  maintained.  The  system,  therefore, 
quickly  reaches  a  new  state  of  equilibrium  in  which  a  higher  vapor 
pressure  exists.  A  heap  of  matter  on  a  table  is  not  in  equilibrium, 
because  addition  of  more  material  produces  no  response  until,  when 
a  very  great  quantity  is  added,  the  table  breaks.  But  a  body  on 
the  scales  is  in  equilibrium,  for  the  addition  of  the  smallest  particle 
produces  a  corresponding  inclination  of  the  beam. 

In  the  preceding  illustration,  the  evaporating  tendency  was 
favored  by  a  rise  in  temperature.  As  an  example  of  a  change  in 
conditions  disfavoring  one  tendency,  take  the  case  where  the  liquid 
is  placed  in  an  open,  shallow  vessel.  Here  the  condensing  tendency 
is  markedly  discouraged,  for  there  is  practically  no  return  of  the 
emitted  molecules.  Hence  complete  evaporation  takes  place.  Ele- 
vation of  the  temperature  hastens  the  process.  A  draft  insures  the 
practical  prevention  of  all  returns,  and  has  therefore  the  same  effect. 
The  two  methods  of  assisting  the  displacement  of  an  equilibrium, 
and  particularly  the  second,  in  which  the  opposed  process  is  weak- 
ened and  the  forward  process  triumphs  solely  on  this  account,  should 
be  noted  carefully.  They  are  applied  with  surprising  effectiveness 
in  the  explanation  of  chemical  phenomena  (see  Chaps.  XV  and 
XVIII). 

Molecular  View  Applied  to  Solids.  —  The  properties  of  solids 
differ  from  those  of  liquids  chiefly  in  the  fact  that  the  solid  has  a 
definite  form  of  which  it  can  be  deprived  only  with  difficulty.  This 
we  may  explain  in  accordance  with  the  kinetic  theory  by  the  supposi- 
tion that  the  cohesion  in  solids  is  very  much  more  prominent  than  in 
liquids.  We  obtain  solids  from  liquids  by  cooling  them;  in  other 
words,  by  diminishing  the  kinetic  energy  and  therefore  the  velocity 
of  the  particles.  The  cohesive  tendency  of  the  latter  is  thus  able  to 
make  itself  felt  to  a  greater  extent.  If,  conversely,  we  heat  a  solid,  or, 
according  to  the  hypothesis,  if  we  increase  the  speed  with  which  the 
particles  move,  the  body  first  melts  and  gives  a  liquid,  and  this  finally 


THE  KINETIC-MOLECULAR  VIEWPOINT  171 

boils  and  becomes  a  gas.  The  intrinsic  cohesion  of  the  particular  sub- 
stance can  undergo  no  change,  but  the  increasing  kinetic  energy  of  the 
particles  steadily  and  continuously  obliterates  its  effects.  Yet  some 
motion  still  survives  in  a  solid.  Thus  we  find  that  when  the  layer  of 
silver  is  stripped  from  a  very  old  piece  of  electroplate  the  presence  of 
this  metal  in  the  German  silver  or  copper  basis  of  the  article  is  easily 
demonstrated. 

Roberts  Austen  has  found  that  if  bars  of  lead  are  prepared,  in  one  end  of  which 
an  alloy  containing  a  certain  proportion  of  gold  has  been  used,  while  the  remainder 
of  the  bar  is  composed  of  pure  lead,  the  gold  has  a  tendency  to  wander  slowly  into 
the  pure  lead.  The  process  is  greatly  aided  by  keeping  the  bars  at  a  fairly  high 
temperature,  but  one  much  below  the  melting-point  is  amply  sufficient.  After  a 
suitable  interval  of  time  the  bar  may  be  sawn  into  fragments  of  equal  length,  and 
its  parts  analyzed.  The  quantity  of  gold  in  a  section  is  found  to  increase  as  we 
approach  the  portion  of  the  bar  to  which  originally  the  whole  of  the  gold  was 
confined. 

The  tendency  of  all  solids  to  assume  crystalline  forms,  which 
show  definite  cleavage  and  other  evidences  of  structure,  distinguishes 
them  sharply  from  liquids  (see  Crystal  structure).  The  force  of 
cohesion  in  liquids  is  exercised  equally  in  different  directions.  In 
solids  it  must  differ  in  different  directions  in  order  that  structure  may 
result.  Since  each  substance  shows  an  individual  structure  of  its 
own,  these  directive  forces  must  have  special  values  in  magnitude  and 
direction  in  each  substance. 

Crystallization.  —  A  crystal  arises  by  growth.  When  the  process 
is  watched,  as  it  occurs  in  a  melted  solid  or  an  evaporating  solution, 
the  slow  and  systematic  addition  of  the  material  in  lines  and  layers, 
as  if  according  to  a  regular  design,  is  one  of  the  most  beautiful  and 
interesting  of  natural  phenomena.  The  fern-like  patterns  produced 
by  ice  on  a  window-pane  show  the  general  appearance  characteristic 
of  crystallization  in  a  thin  layer.  A  larger  mass  in  a  deep  vessel  gives 
forms  which  are  geometrically  more  perfect.  From  its  very  incipiency 
the  crystal  has  the  same  form  as  when,  later,  its  outlines  can  be  dis- 
tinguished by  the  eye.  Hence  the  outward  form  is  only  an  expression 
of  a  specific  internal  structure  which  the  continual  reproduction  of 
the  same  outward  form  on  a  larger  and  larger  scale  leaves  as  a  me- 
morial of  itself  in  the  interior. 

Crystal  Forms.  —  Crystalline  form  is  so  continually  used  in 
identifying  (p.  40)  the  substances  produced  in  chemical  actions  that 


172 


INORGANIC  CHEMISTRY 


a  list  of  the  kinds  of  forms  which  occur  will  assist  in  giving  definite 
meaning  to  our  descriptions. 

The  classification  of  crystalline  forms  is  carried  out  according  to 
the  degree  of  symmetry  of  the  crystals.  Thirty-two  distinct  classes 
are  distinguished,  but  for  our  purpose  a  rougher  division  into  six 
groups  will  suffice.  These  groups  are  known  by  the  following  names : 

1.  Regular  system.  5.  Monosymmetric,  or  monoclinic 

2.  Square  prismatic  system.  system. 

3.  Hexagonal  system.  6.  Asymmetric,  or  triclinic  system. 

4.  Rhombic  system. 

The  regular  system  presents  the  most  symmetrical  figures  of  all. 
Some  forms  which  commonly  occur  are  the  octahedron  (Fig.  63) 
shown  by  alum,  the  cube  (Fig.  64)  affected  by  common  salt,  and 
the  dodecahedron  (Fig.  65)  frequently  assumed  by  the  garnet. 


FIG.  63. 


04. 


FIG.  65. 


The  square  prismatic  system  includes  less  symmetrical  forms  than 
the  previous  one,  since  the  crystals  are  lengthened  in  one  direction. 
Fig.  66  shows  the  condition  in  which  zircon  ZrSiO4,  which  furnishes 
us  with  the  basis  of  certain  incandescent  illuminating  arrangements, 
occurs  in  nature.  The  form  of  ordinary  hydrated  nickel  sulphate 
NiSO4;6H2O  is  similar  to  this. 


FIG.  67. 


FIG.  68. 


The  hexagonal  system,  like  the  preceding,  frequently  exhibits 
elongated  prismatic  forms,  but  the  section  of  the  crystals  is  a  hexagon 
instead  of  a  square,  and  the  termination  is  a  six-sided  pyramid. 
Quartz  (Fig.  67),  or  rock  crystal,  is  the  most  familiar  mineral  in  this 
system.  Calcite  (CaCO3),  which  is  chemically  identical  with  chalk, 


THE  KINETIC-MOLECULAR  VIEWPOINT 


173 


or  marble,  takes  forms  known  as  the  scalenohedron  (Fig.  68)  and 
rhombohedron  (Fig.  69),  which  are  classified  in  a  subdivision  of  this 
system.  Indeed,  recently  it  has  become  common  to  erect  this  into 
a  separate  system  (the  trigonal),  in  which  both  quartz  and  calcite  are 
included. 

The  rhombic  system  includes  the  natural  forms  of  the  topaz,  and 
of  sulphur  (Fig.  8,  p.  14),  as  well  as  those  of  potassium  permanganate 
(Fig.  70),  potassium  nitrate  (Fig.  71),  and  many  other  substances. 
These  crystals  exhibit  a  good  deal  of  symmetry,  but  their  section  is 
always  rhombic,  and  hence  the  name. 


FIG.  09. 


FIG.  70. 


FIG.  71. 


The  monosymmetric  system  exhibits  forms  which  have  but  one 
plane  of  symmetry.  Gypsum  (Fig.  72),  which  is  hydrated  calcium 
sulphate  (CaS04,2H20),  and  felspar  (Fig.  3,  p.  5)  are  minerals 
possessing  forms  of  this  kind.  Tartaric  acid,  rock  candy  (Fig.  73), 
potassium  chlorate,  and  hydrated  sodium  carbonate  (washing  soda) 
belong  to  this  system. 


FIG.  72. 


FIG.  73. 


The  asymmetric  system  includes  forms  which  have  no  plane  of 
symmetry  whatever.  Blue  vitriol  (Fig.  55,  p.  151),  CuSO4,5H20,  is 
one  of  the  most  familiar  substances  of  this  kind. 

The  forms  of  crystals  which  we  may  actually  make  seldom  corre- 
spond exactly  with  the  figures.  If  we  allow  a  crystal  to  grow  upon 
the  bottom  of  a  vessel,  for  example,  it  will  usually  have  a  tendency  to 
spread  itself  out  parallel  to  the  surface  of  the  glass,  and  when  taken 
up  for  examination,  will  be  found  to  present  a  somewhat  distorted 
form.  By  changing,  at  frequent  intervals,  the  face  on  which  the 


174  INORGANIC  CHEMISTRY 

crystal  stands,  however,  uniform  growth  in  all  directions  is  secured. 
Hanging  a  small  crystal  by  a  thread  insures  almost  ideal  development. 
Yet  the  form  even  of  distorted  crystals  can  readily  be  recognized  by 
suitable  means.  The  shape  of  the  faces  may  indeed  be  extremely 
misleading.  We  find,  however,  that  the  angles  at  which  the  faces 
meet  are  always  the  same,  whatever  disproportionate  growth  may 
have  occurred  in  the  development  of  the  crystal. 

Since,  in  general,  each  substance  has  a  form  of  its  own,  no  other 
substance,  as  a  rule,  can  be  used  even  partially  in  building  up  the 
crystal  (see,  however,  Isomorphism).  This  fact  is  taken  advantage 
of  in  order  to  separate  chemical  substances  from  impurities.  The 
impure  body  is  first  dissolved  in  some  solvent.  The  preponderating 
substance  in  the  mixture,  unlass  it  is  very  much  more  soluble  than 
the  impurity,  will  then  usually  give  pure  crystals  while  the  foreign 
body  remains  in  solution. 

The  shapes  of  gems  must  not  be  confused  with  crystalline  forms. 
The  original  crystals  are  cut  and  polished  to  a  new  form  specially 
adapted  to  increase  the  ornamental  value  of  the  stone  by  causing  it  to 
reflect  more  light  (see  Diamond) .  Again,  glass  is  really  a  very  viscous 
fluid  (amorphous  body,  p.  154),  and  has  no  structure  or  form  of  its 
own.  The  word  " crystal"  applied  to  cut  glass  would  therefore  be 
misleading  if  taken  literally. 

Crystal  Structure.  —  As  the  above  would  lead  us  to  expect,  the  study  of 
crystallized  substances  shows  that  their  peculiarities  are  not  confined  to  the  out- 
side layer.  The  outline  represents  a  certain  structure  which  permeates  the  whole 
mass.  In  crystals  of  the  regular  system,  many  of  the  ordinary  physical  properties 
are  the  same  as  those  of  an  amorphous  substance,  like  glass.  For  example,  if  we 
turn  a  sphere  out  of  crystallized  salt  and  hang  it  in  pure  water,  we  find  that  solu- 
tion takes  place  at  a  uniform  rate  all  over  the  surface.  This  is  not  the  case,  how- 
ever, with  substances  from  any  of  the  other  systems.  Spheres  cut  from  substances 
belonging  to  the  second  and  third  systems  would  dissolve  more,  or  less,  rapidly  in 
the  direction  of  the  chief  axis  than  in  any  other  direction,  and  so  ellipsoids  of 
revolution  would  quickly  be  produced.  In  the  other  three  systems,  more  complex 
forms  would  result. 

The  tenacity  of  crystals,  to  whatever  system  they  may  belong,  is  different  in 
different  directions.  Thus,  a  crystal  of  salt  has  a  cleavage  parallel  to  any  of  the 
faces  of  the  cube,  and,  therefore,  splits  most  easily  in  one  of  three  directions  at 
right  angles  to  each  other.  Calcite,  whatever  the  outward  form  of  the  crystal, 
always  cleaves  so  as  to  give  a  rhombohedron.  Fluorite  CaF2,  although  almost 
always  cubical  in  form,  splits  when  broken  so  as  to  give  an  octahedron. 

The  behavior  of  crystals  towards  light  is  also  extremely  interesting.  The  rate 
at  which  light  moves  through  crystals  of  the  regular  system  is  the  same  in  all 


THE  KINETIC-MOLECULAR  VIEWPOINT  175 

directions.  In  other  crystals,  however,  we  find  that  it  moves  with  a  different 
speed  in  different  directions,  the  variations  in  speed  being  likewise  related  to  the 
outward  form.  Finally,  if  a  thin  slab  of  rock  salt  is  covered  with  wax  and  the 
point  of  a  heated  cone  of  metal  is  placed  in  the  center,  the  wax  melts  uniformly' 
in  a  circle  around  the  point,  indicating  that  the  heat  is  conducted  with  equal  speed 
in  all  directions.  The  way  in  which  the  slab  has  been  cut  with  reference  to  the 
surface  of  the  crystal,  in  the  case  of  substances  of  the  first  of  the  above  systems, 
has  no  effect  upon  this  result.  In  all  other  cases,  however,  the  zone  of  melting 
wax  is  in  general  elliptical,  or  even  more  complex  in  form,  according  to  the  system 
to  which  the  substance  belongs  and  the  direction  in  which  the  slab  has  been  cut 
((see  Chap.  XXII). 

Practical  Value  of  the  Molecular  Viewpoint.  —  The  value 
of  the  kinetic-molecular  viewpoint  has  been  illustrated  above,  and 
will  appear  again  when  we  deal  with  solutions  and  chemical  equi- 
librium. But  indeed  its  field  of  application  is  coextensive  with  the 
science  itself. 

The  conceptions  of  the  kinetic  theory  are  of  especially  great  assist- 
ance in  rationalizing  chemical  manipulation  and  so  hastening  the 
acquisition  of  an  intelligent  control  of  it.  The  constructive  imagina- 
tion, on  the  use  of  which  experimental  work  depends  so  much  for  its 
success,  must  have  something  to  work  with,  and  this  view  furnishes 
a  tool  such  as  it  requires.  Methods  taught  by  rule  of  thumb  are 
slowly  learned  and  constantly  fail  in  application.  We  may  be  told 
a  dozen  times  that  using  reagents  in  finely  powdered,  or  metals  in 
granulated  condition  hastens  all  interactions,  and  still  never  think  of 
this  abstraction  when  working.  But  if  it  is  suggested  that,  in  terms 
of  the  kinetic  theory,  molecules  must  meet  freely  in  the  same  medium 
to  react  easily,  and  that,  therefore,  the  larger  the  surface  the  more 
copious  will  be  the  supply  of  molecules  dissolving,  we  are  likely  to 
form  a  conception  of  the  reason  for  the  procedure  that  will  be 
lasting. 

When  a  student  is  told  to  concentrate  a  solution  and  set  it  aside 
to  crystallize,  why  does  he  evaporate  the  liquid  rapidly  to  dryness, 
and  still  expect  the  residue  to  appear  in  large,  well-formed  crystals? 
Because  he  has  no  notion  of  the  necessary  slowness  of  the  process  of 
crystallization.  But  if  he  has  the  idea  that  it  is  like  building  a  house, 
one  stone  at  a  time,  and  that  there  are  far  more  units  to  be  laid  down 
according  to  plan  in  making  the  smallest  visible  crystal  than  there  are 
bricks  in  building  the  largest  factory,  his  procedure  may  promptly 
become  more  rational. 


176  INORGANIC  CHEMISTRY 

Stochastic  Hypotheses.  —  As  we  have  seen  (p.  11),  one  kind 
of  explanation  consists  in  imagining  some  fact  or  connection  of  facts, 
or  some  machinery  which,  if  it  existed,  would  account  for  the  phe- 
nomena to  be  explained.  This  we  call  an  hypothesis.  Hypotheses 
are  of  two  kinds,  stochastic  and  formulative. 

The  first  of  these  is  constantly  used  at  every  step  in  investigation, 
although  it  is  seldom  mentioned  in  books.  When  Mitscherlich 
discovered  that  Glauber's  salt  (p.  151)  gave  a  definite  pressure  of 
water  vapor,  he  at  once  formed  the  hypothesis,  that  is,  supposition, 
that  other  hydrates  would  be  found  to  do  likewise.  Experiments 
showed  this  supposition  to  be  correct.  The  hypothesis  was  at  once 
displaced  by  the  fact.  This  sort  of  hypothesis  predicts  the  probable 
existence  of  certain  facts  or  connections  of  facts,  hence,  reviving  a 
disused  word,  we  call  it  a  stochastic  hypothesis  (Gk.  o-roxao-riKos,  apt 
to  divine  the  truth  by  conjecture).  It  professes  to  be  composed 
entirely  of  verifiable  facts  and  is  subjected  to  verification  as  quickly 
as  possible. 

Formulative  Hypotheses.  —  The  formulative  hypothesis  is 
mainly  a  structure  existing  in  the  imagination.  It  may  not  be 
capable  of  existing  anywhere  else,  because  it  may  include  items  which 
are  more  or  less  in  conflict  with  ordinary  experience.  But  then,  we 
use  a  formulative  hypothesis  for  the  purpose  of  having  before  our 
minds  some  sort  of  machinery  that  will  help  us  to  understand  some 
aspect  of  the  behavior  of  matter  or  energy,  and  we  make  few,  if  any, 
attempts  to  verify  the  hypothesis  itself.  Thus,  the  conception  of 
light  as  consisting  of  corpuscles  was  such  an  hypothesis  in  its  day. 
The  imponderable-matter  (note  the  contradiction)  view  of  heat  was 
another.  The  undulatory  theory  of  light,  which  postulates  a  per- 
fectly elastic  ether,  weightless,  frictionless,  and  lacking  every  trace  of 
impenetrability,  is  an  hypothesis  showing  the  same  inverifiable 
elements.  It  is  handled  characteristically  also,  for  we  do  not  try  by 
its  means  to  learn  more  about  ether,  but  more  about  light. 

Such  an  hypothesis  is  like  a  scaffolding,  constructed  to  enable  us 
to  examine  or  work  upon  a  difficultly  accessible  part  of  a  building, 
which  we  never  for  a  moment  think  of  as  being  a  part  of  the  building. 
It  is  a  sort  of  formula.  The  algebraic  formula  represents  magni- 
tudes; the  geometrical,  directions  and  dimensions.  The  formula  in 
physics,  by  the  use  of  mathematical  conventions,  pictures,  for 
example,  some  mode  of  behavior  of  matter.  For  so  concrete  a  subject, 
however,  the  mathematical  mode  of  expression  is  intensely  abstract. 


THE  KINETIC-MOLECULAR  VIEWPOINT  177 

And  so  a  representation  in  terms  of  mechanism,  which  is  still  a 
formula,  is  frequently  resorted  to.  The  hypothesis  of  the  ether  is 
therefore  a  formula  consisting  of  imaginary  machinery.  Its  object 
is  simply  to  help  us  in  organizing  or  formulating  knowledge  on  a 
certain  subject.  Hence  we  name  it  a  formulative  hypothesis. 

The  kinetic  molecular  view  was  at  first  an  hypothesis  of  this  kind. 
It  is  only  recently  that  we  have  been  able  to  count  the  molecules  of 
helium  given  off  by  radium  (q.v.)  and  to  see  and  to  photograph  the 
path  of  each  molecule.  These,  and  other  observations,  have  now 
shown  that  molecules  really  exist,  and  behave  as  the  hypothesis 
assumed  they  did.  But  the  conception  of  molecules  was  no  less 
useful  before  it  had  been  verified  than  it  was  afterwards.  Dalton, 
who  first  worked  out  a  clear  conception  of  the  independent  behavior 
of  mixed  gases  (p.  Ill),  used  this  as  an  hypothesis  constantly  in  that 
work.  Clear  understanding  of  the  separate  existence  of  aqueous 
vapor  in  the  air  could  be  reached  by  him  only  by  thinking  of  each 
material  as  being  made  up  of  independent  molecules.  Any  other 
mode  of  conceiving  the  mixture  that  might  readily  occur  to  one  would 
involve  some  adhesion  or  interference  of  the  two  substances.  His 
recognition  of  the  independent  solubilities  of  mixed  gases,  in  propor- 
tion to  the  partial  pressures  of  each,  would  have  been  delayed  or 
prevented  altogether  if,  in  studying  the  results  of  his  experiments,  he 
had  not  reached  it  by  way  of  this  hypothesis. 

Exercises.  —  1.  What  are  the  relative  rates  of  diffusion  of  am- 
monia gas  (density  17)  and  of  sulphur  dioxide  (density  64)? 

2.  What  are  the  speeds  in  meters  per  second  of  the  molecules  of 
ammonia  and  of  sulphur  dioxide  at  room  temperature? 

3.  Give  two  ways  of  obtaining  crystals  of  a  substance. 


CHAPTER  X 
SOLUTION 

WE  have  frequently  made  use  of  the  fact  that  certain  substances 
form  with  others  homogeneous  systems  which  we  call  solutions. 
Sometimes  this  property  is  taken  advantage  of  for  separating  ma- 
terials, as  in  the  case  of  the  removal  of  sulphur  from  admixture  with 
iron  or  ferrous  sulphide  (p.  15).  In  other  cases  we  carry  out  the 
interaction  of  chemical  substances,  by  first  dissolving  each  in  some 
liquid  and  then  mixing  the  solutions.  The  liquid,  commonly  water,  is 
used  as  a  vehicle  for  one  or  more  of  the  substances,  and  often  takes  no 
part  in  the  chemical  change.  Thus  some  knowledge  of  the  properties 
of  solutions  is  absolutely  necessary  in  order  that  we  may  employ  them 
intelligently.  In  what  follows,  we  shall  give  a  preliminary  account  of 
some  of  the  simpler  facts  about  solution. 

General  Properties  of  Solutions.  —  A  solid  may  be  distrib- 
uted through  a  liquid  either  by  being  simply  suspended  (p.  142)  in  the 
latter  (mixture)  or  by  being  dissolved  in  it  (solution).  Similarly,  a 
liquid  may  be  suspended  in  droplets  in  another,  like  the  fat  in  milk 
(emulsion),  or  it  may  be  dissolved.  It  is  usually  easy  to  distinguish 
between  the  two  cases,  for  a  suspiended  substance  settles  or  separates 
sooner  or  later  (like  the  fats  in  milk  —  as  cream)  while  a  dissolved  one 
shows  no  such  tendency.  The  cases  are  exceptional  where  the  sub- 
division of  a  suspended  substance  is  so  iriinute  (colloidal  suspension, 
q.v.)  as  tc  make  its  retention  by  filter  paper  impossible.  If  a  liquid  is 
opalescent  or  opaque,  then  we  have  a  case  of  suspension.  A  solution 
is  a  clear,  transparent,  perfectly  homogeneous  liquid,  in  which  the 
dissolved  substance  seems  to  have  been  dispersed  so  completely  that 
the  liquid  cannot  be  distinguished  by  the  eye  from  a  pure  substance. 

There  is  no  limit  to  the  amount  of  dissipation  which  may  thus  be 
produced.  A  single  fragment  of  potassium  permanganate,  for  ex- 
ample, which  gives  a  very  deep  purple  solution  in  water,  may  be  dis- 
solved in  a  liter  or  even  in  twenty  liters  of  water,  and  the  purple 
tinge  which  it  gives  to  the  liquid  will  still  be  perfectly  perceptible  in 
every  part  of  the  larger  volume.  The  qualitative  characteristics, 

178 


SOLUTION  179 

therefore,  of  solution  are  absence  of  settling,  homogeneity,  and  ex- 
tremely minute  subdivision  of  the  dissolved  substance. 

The  Scope  of  the  Word.  —  The  word  is  used  for  other  systems 
than  those  containing  a  solid  body  dissolved  in  a  liquid.  Thus,  liquids 
also  may  be  dissolved  in  liquids,  as  alcohol  in  water.  Again,  if  we 
warm  ordinary  water,  bubbles  of  gas  appear  on  the  sides  of  the  vessel 
before  the  water  has  approached  the  boiling-point.  They  are  found 
to  be  gas  derived  from  the  air.  Agitation  of  any  gas  with  water 
results  in  the  solution  of  a  large  or  small  quantity  of  the  gas,  and 
heating  will  usually  drive  the  gas  out  again.  It  appears  therefore 
that  solids,  liquids,  and  gases  can  equally  form  solutions  in  liquids. 

The  absorption  of  hydrogen  by  palladium  (at  all  events  after  a 
certain  point)  and  by  iron  takes  place  in  accordance  with  the  same 
laws  as  the  solution  of  solids  in  liquids,  and  the  results  may  be  de- 
scribed therefore  as  true  solutions.  Liquids  are  in  some  cases 
absorbed  by  solids,  and  homogeneous  mixtures  of  solids  with  solids 
are  perfectly  familiar.  The  sapphire  is  a  solution  of  a  small  amount 
of  a  strongly  colored  substance,  in  a  large  amount  of  colorless  alumin- 
ium oxide.  It  may  therefore  be  stated  that  solution  of  gases,  liquids, 
and  solids  in  solids  appears  to  be  possible. 

Limits  of  Solubility.  —  The  next  question  which  naturally 
occurs  to  us  is  as  to  whether  the  mingling  of  two  substances  in  this 
manner  has  any  limits.  We  find  that  the  results  of  experiment  in  this 
direction  may  be  divided  into  two  classes.  Some  pairs,  of  liquids 
particularly,  may  be  mixed  in  any  proportions  whatever.  Alcohol 
and  water,  and  glycerine  and  water  are  such  pairs.  On  the  other 
hand,  at  the  ordinary  laboratory  temperature,  we  can  scarcely  take 
a  fragment  of  marble  Ca©O3  so  small  that  it  will  dissolve  completely 
in  100  c.c.  of  pure  water,  for  only  0.00013  g.  dissolves.  Under  the 
same  conditions  any  amount  of  potassium  chlorate  up  to  about  5  g. 
will  completely  disappear  after  vigorous  stirring,  while  90  g.  of 
ordinary  Epsom  salts  (hydrated  magnesium  sulphate),  but  not  more, 
may  be  dissolved  in  about  the  same  amount  of  water.  In  fact,  most 
solids  may  be  dissolved  in  a  liquid  only  up  to  a  certain  limit,  which 
with  different  solids  may  range  from  a  scarcely  perceptible  up  to  a 
very  large  amount.  No  substance  is  absolutely  insoluble.  But  for 
the  sake  of  brevity  we  call  marble,  for  example,  insoluble,  because  in 
most  connections  it  may  be  so  considered. 

Chemists  have  not  yet  succeeded  in  explaining  these  differences 
in  solubility,  which  are  often  so  surprising.  Thus,  guncotton  (cellu- 


180 


INORGANIC   CHEMISTRY 


lose  nitrate)  is  soluble  in  a  mixture  of  alcohol  and  ether,  but  not  in 
these  liquids  separately.  On  the  other  hand,  cellulose  acetate,  an 
allied  substance,  used  in  making  bristles  for  brushes  and  artificial 
horsehair,  is  soluble  in  these  liquids  separately,  but  not  in  the  mix- 
ture! Again,  silver  fluoride  is  extremely  soluble  in  water,  while  the 
allied  chloride,  bromide,  and  iodide  of  silver 
are  very  insoluble  (see  Table,  inside  front 
cover).  Conversely,  calcium  fluoride  is  very 
insoluble,  but  the  chloride,  bromide,  and  iodide 
are  quite  soluble. 

Recognition  and  Measurement  of  Sol- 
ubility. —  The   only   method  of   recognizing 

with  certainty  whether  a  solid  is  soluble  in  a 
liquid  or  not  is  to  filter  the  mixture  and  evap- 
orate a  few  drops  of  the  filtrate  on  a  clean 
watch-glass.  For  learning  how  much  of  the 
body  is  contained  in  a  given  solution,  a  weighed 
quantity  of  the  solution  is  evaporated  to  dry- 
ness  and  the  weight  of  the  residue  determined. 
When  the  dissolved  substance  is  volatile,  its 
presence  is  often  shown  by  some  chemical  test 
(p.  89). 

Ether  and  water  is  a  case  typical  of  the 
behavior  of  two  liquids,  each  somewhat  soluble 
in  the  other.  After  being  shaken  together,  they  seem  to  separate 
again  completely  into  two  layers  (Fig.  74),  with  the  ether  uppermost. 
If,  however,  the  water  is  withdrawn  from  beneath  the  ether,  we  find 
that,  when  heated,  it  gives  off  quantities  of  ether  vapor  which  can 
be  set  on  fire.  Conversely,  the  addition  of  anhydrous  cupric  sulphate 
to  a  sample  of  the  ether  shows  the  presence  of  water  in  the  latter,  for 
the  blue  hydrated  form  of  the  substance  is  at  once  produced.  In 
some  common  cases  the  maximum  solubilities  at  room  temperature 
are  as  follows: 


FIG.  74. 


SUBSTANCE. 

VOL.  IN 
100  VOLS.  AQ. 

VOL.  AQ.  IN 
100  VOLS. 

GRAMS 
IN  100  G.  AQ. 

GRAMS  AQ. 

IN   100   G. 

Alcohol          

No  limit 

No  limit 

No  limit 

No  limit 

Ether         

8  11 

2  93 

5  97 

3  98 

Chloroform       

0.42 

0.15 

0  66 

0.10 

Carbon  disulphide  .    .    . 

0.17 

0.96 

0.22 

0.74 

SOLUTION  181 

It  must  be  stated  explicitly  that  in  going  into  solution,  as  we  have 
used  the  term,  a  compound  dissolves  as  a  whole  and,  if  the  compound 
is  pure  (p.  6),  any  residue  has  the  same  chemical  composition  as  the 
part  which  has  dissolved.  If  the  undissolved  residue  is  a  different 
substance,  a  chemical  interaction  with  the  solvent  has  occurred.  If, 
on  evaporation,  a  different  substance  remains,  there  has  also  been 
chemical  action. 

Terminology.  —  In  order  to  describe  the  relations  of  the  com- 
ponents of  a  solution,  certain  conceptions  and  corresponding  technical 
expressions  are  required. 

It  is  customary  to  speak  of  the  substance  which,  like  water  in 
most  cases,  forms  the  bulk  of  the  solution,  as  the  solvent.  To  express 
the  substance  which  is  dissolved,  the  word  solute  is  frequently  used, 
and  will  be  employed  when  we  wish  to  avoid  circumlocution. 

The  term  " strength"  is  too  indefinite  for  scientific  purposes.  It 
may  imply  activity,  or  power  of  resistance,  or  pungency  (in  an  odor), 
or,  as  in  the  case  of  solutions,  it  may  be  a  measure  of  quantity.  The 
amount  of  the  substance  which  has  been  dissolved  by  a  given  quantity 
of  the  solvent  is  therefore  described  as  the  concentration  of  the  solu- 
tion. A  solution  containing  a  small  proportion  of  the  dissolved  body 
is  called  dilute;  it  has  a  small  concentration.  One  which  contains  a 
larger  amount  is  more  concentrated.  Very  " strong"  solutions  are 
frequently  spoken  of  simply  as  concentrated  solutions.  The  partial 
removal  of  the  solvent  by  evaporation  is  called  concentrating,  its  total 
removal  evaporating  to  dryness. 

Finally,  since  there  is  a  limit  to  the  solubility  of  most  substances, 
a  solution  is  described  as  saturated  when  the  solute  has  given  as  much 
material  to  the  solvent  as  it  can.  This  state  is  reached  after  pro- 
longed agitation  with  an  excess  of  the  gas,  of  the  liquid,  or  of  the  finely 
powdered  solid,  as  the  case  may  be  (see  pp.  192,  194).  The  larger  the 
excess,  the  sooner  saturation  is  attained,  but  the  concentration  is  no 
greater.  The  maximum  concentration  attainable  in  this  way  is 
called  the  solubility  of  the  substance  in  a  given  solvent.  Note  that 
a  saturated  solution  need  not  also  be  a  concentrated  one.  It  will  be 
very  dilute,  if  the  solute  is  but  slightly  soluble. 

The  distinction  between  solute  and  solvent  is  made  merely  for 
convenience.  Theoretically,  there  is  no  distinction  between  the 
components  of  a  solution. 

Units  Used  in  Expressing  Concentrations.  —  The  concen- 
trations of  solutions,  saturated  and  otherwise,  are  sometimes  ex- 


182 


INORGANIC  CHEMISTRY 


pressed  in  physical,  and  sometimes  in  chemical,  units  of  weight. 
When  physical  units  are  employed,  as  in  the  above  table,  we  give 
the  number  of  grams  of  the  solute  held  in  solution  by  one  hundred 
grams  (in  the  case  of  water  100  g.  =  100  c.c.)  of  the  solvent. 

The  solubilities  at  18°  of  142  bases  and  salts  are  given  in  a  table, 
placed  inside  the  front  cover  of  this  book. 

When  chemical  units  of  weight  are  employed,  two  different  plans 
are  possible,  and  both  are  in  use.  Either  the  equivalent  (p.  63)  or  the 
molecular  weights  may  be  taken  as  a  basis  of  measurement.  In  the 
former  case,  the  solutions  are  called  normal  solutions,  and  in  the  latter, 
for  a  reason  which  will  appear  later  (Chap.  XII),  molar  solutions. 

A  normal  solution  contains  one  gram-equivalent  of  the  solute  in  one 
liter  of  solution  (not  in  1  1.  of  solvent).  The  word  " equivalent" 
has  been  used  hitherto  only  of  elements,  and  this  application  of  the 
expression  involves  an  extension  of  its  meaning.  An  equivalent 
weight  of  a  compound  is  that  amount  of  it  which  will  interact  with 
one  equivalent  of  an  element.  Thus,  a  formula-weight  of  hydro- 
chloric acid  HC1  (36.5  g.)  is  also  an  equivalent  weight,  for  it  contains 
1  g.  of  hydrogen,  and  this  amount  of  hydrogen  is  displaceable  by  one 
equivalent  weight  of  a  metal.  A  formula-weight  of  sulphuric  acid 
H2S04  (98  g.),  however,  contains  two  equivalents  of  the  compound, 
and  a  formula-weight  of  aluminium  chloride  A1C13  (133  5  g.)  three 
equivalents.  Hence  normal  solutions  of  these  three  substances 
contain  respectively  36.5  g.  of  HC1,  \  of  98  g.,  or  49  g.  of  H2SO4,  and 
f  of  133.5  g.,  or  44.5  g.  of  A1C13  per  liter  of  solution.  The  special 
property  of  normal  solutions  is,  obviously,  that  equal  volumes  of  two 
of  them  contain  the  exact  proportions  of  the  solutes  which  are  re- 
quired for  complete  interaction.  Solutions  of  this  kind  are  much 
used  in  quantitative  analysis. 

Solutions  of  different  concentrations  all  prepared  on  the  above 
basis  are  named  as  follows,  and  are  often  indicated  by  the  abbrevia- 
tions appended: 


QUANTITY  OF  SOLUTE  PER  LITER. 

NAME. 

ABBREVIATIONS. 

One  hundredth  of  one  gram-  ) 
equivalent      ) 

Centi-normal 

^  or  .01  N 

One  tenth  of  one  gram-equiv-  ) 
alent     \ 

Deci-normal 

16  OT-IN 

One  half  of  one  gram-equiva-  ) 
lent    } 

Semi-normal 

1U 

N 
-z     OT.5N 

One  gram-equivalent      .... 
Two  and  a  half  gram-equivalents 

Normal 
Two  and  a  half  normal 

A 

N 
2%Nor2.5N 

SOLUTION  183 

A  molar  solution  contains  one  mole  (gram-molecular  weight,  see 
p.  183)  of  the  solute  in  one  liter  of  solution  (not  in  1  1.  of  solvent).  When 
molecular  formulae  (see  Chap.  XII)  are  used,  this  means  one  gram-formula  weight 
per  liter.  In  the  cases  cited  above,  the  molar  solution  contains  36.5  g.  HC1, 
98  g.  H2SO4,  and  133.5  g.  A1C13  per  liter.  As  will  be  seen,  the  concentration 
of  molar  and  normal  solutions  are  necessarily  identical  when  the  radicals  are 
univalent.  Other  concentrations  are  described  as  deci-molar  (M/10  or  .1  M), 
two  and  a  half  molar  (2.5  M),  and  so  forth,  on  the  same  plan  as  before. 

There  is  also  a  chemical  unit  of  volume  (see  Chap.  XII)  which  is  the  volume 
occupied  by  a  mole  (gram-molecular  weight)  of  a  gas  (or  dissolved  substance)  at 
0°  and  760  mm.  pressure  (gaseous  or  osmotic).  This  volume  averages  22.4  liters, 
and  is  called  the  gram-molecular  or  molar  volume  (G.M.V.) .  The  unit  of  concen- 
tration for  many  theoretical  purposes  is,  therefore,  that  of  one  mole  in  22.4  liters. 

Solution  One  of  the  Physical  States  of  Aggregation  of 
Matter.  —  When  a  solid  body  dissolves  in  a  liquid,  the  properties 
of  the  body  undergo  a  very  marked  change,  which  to  all  appearance 
might  be  chemical.  Yet,  besides  the  ease  with  which  the  liquid  may 
be  removed  by  evaporation  and  the  solid  recovered  unchanged,  we 
note  particularly  that  the  concentration  of  a  saturated  solution  cannot 
be  expressed  in  terms  of  integral  multiples  of  the  atomic  weights.  If, 
therefore,  the  process  were  to  be  regarded  as  chemical,  several  impor- 
tant generalizations  would  have  to  be  revised  or  discarded  (cf.  p.  61). 
We  shall  see  also,  in  a  later  paragraph,  that  the  quantity  of  a  solid 
which  a  liquid  may  take  up  varies  with  the. slightest  change  in  tem- 
perature. Now  we  do  not  find  the  composition  of  chemical  com- 
pounds so  to  vary.  The  solution  of  a  solid  may  therefore  be  likened 
to  a  chang ;  in  state,  similar  to  the  conversion  of  a  liquid  into  a  gas  or 
a  solid  (cf.  p.  49). 

As  there  is  danger  of  confusion  arising,  we  may  repeat  that  a  com- 
pound is  homogeneous  and  its  composition  is  expressible  in  integral 
multiples  of  the  chemical  units  of  weight  (atomic  weights) ;  a  satu- 
rated solution  is  homogeneous  but  ifs  concentration  varies  gradually 
with  temperature  so  that  integral  multiples  of  the  chemical  units 
cannot  be  used  to  describe  its  composition;  a  mixture  of  two  solids  or 
an  emulsion  of  two  liquids,  is  neither  homogeneous  nor  in  any  way 
definite  in  composition. 

Kinetic- Molecular  View  of  the  State  of  Solution.  —  Ac- 
cepting solution  as  a  physical  state  of  aggregation,  we  may  now  apply 
the  same  molecular  conceptions  to  the  explanation  of  the  behavior  of  a 
substance  in  solution  as  to  matter  in  the  gaseous  or  liquid  states.  We 


184 


INORGANIC  CHEMISTRY 


saw  that  a  solid  body,  which  is  ordinarily  condensed  in  a  small  space, 
can  be  disseminated  by  the  use  of  a  solvent  through  a  very  large  one. 
The  molecules  of  the  solid  become  scattered,  like  those  of  a  gas  or 
vapor,  through  a  much  greater  volume.  We  may  regard  the  dis- 
solved substance  as  being,  practically,  in  a  gaseous  or  quasi-gaseous 
condition.  The  molecules  are  torn  apart  from  one  another,  their 
cohesion  is  overcome,  and  their  freedom  of  motion  is  in  a  measure 
restored.  It  is  true  that  they  could  not  continue  to  occupy  this  large 
volume  for  a  moment  in  the  absence  of  the  solvent.  But  we  may 
bring  this  into  relation  with  the  case  of  a  vapor  by  saying  that  a  solid 
body,  like  common  salt,  can  only  evaporate  (i.e.,  dissolve)  appreciably 
at  the  ordinary  temperature,  and  occupy  a  large  space,  when  that 
space  is  already  filled  with  a  suitable  liquid.  The  latter  acts  as  a 
vehicle  for  the  particles  of  the  solid.  A  volatile  liquid,  on  the  con- 
trary, can  dissolve  in  an  empty  space  and  fill  it  with  its  particles 
without  any  vehicle  being  required. 


FIG.  75. 


FIG.  76. 


This  conception  of  the  quasi-gaseous  condition  of  a  dissolved  sub- 
stance would  be  simply  fantastic  if  it  did  not  lead  us  to  a  better  under- 
standing of  the  behavior  of  solutions.  It  does  successfully  explain 
many  facts  about  solutions,  such  as  diffusion,  and  saturation  (see 
next  section). 

It  is  easy  to  show  that,  if  we  place  a  quantity  of  the  pure  solvent 


SOLUTION 


185 


(Fig.  75)  above  a  concentrated  solution  of  a  suostance,  and  then  set 
the  arrangement  aside,  the  dissolved  body  slowly  makes  its  way 
through  the  liquid  (Fig.  76),  obliterating  the  original  plane  of  separa- 
tion. Eventually  the  dissolved  body  scatters  itself  uniformly  through 
the  whole.  In  other  words,  the  particles  of  the  dissolved  substance 
exhibit  the  property  of  diffusion  in  the  same  way  as  do  those  of 
gases. 

When  the  diffusion  of  a  gas  is  resisted  by  a  suitable  partition,  we 
find  pressure  is  exercised  upon  the  walls  of  the  vessel  and  upon  the 
partition.  It  is  possible  to  show  that  the  particles  of  a  dissolved  sub- 
stance  induce  a  pressure  of  a  very  similar  kind.  This  pressure  is 
spoken  of  as  diffusion  pressure,  and  is  proportional  to  the  concen- 
tration of  the  solution.  In  other  words,  it  depends  upon  the  degree 
of  crowding  of  the  molecules  of  the  dissolved  substance. 

Kinetic- Molecular  View  of  the  Process  of  Solution.  —  We 

may  now  apply  the  same  ideas  to  the  process  of  dissolving,  with  a 
view  more  especially  to  explaining  why  the  process  of  dissolving 
ceases,  in  spite  of  the  presence  of  excess  of  the 
solute,  when  a  certain  concentration  has  been 
reached.     If  some  of  the  material  dissolves,  why 
not  more? 

Let  us  suppose  that  it  is  the  dissolving  of 
commoA  salt  in  water  (Fig.  77)  which  we  wish  to 
explain  in  detail.  We  believe  that  in  the  solid 
substance  the  molecules  are  closely  packed  to- 
gether, while  in  the  solution  they  are  rather 
sparsely  distributed.  If  there  were  no  water  over 
the  salt,  practically  none  of  the  particles  of  the 
latter  would  be  able  to  leave  the  solid  and  enter 
the  space  above.  Thus,  the  process  of  solution 
must  consist  in  the  loosening  of  the  molecules 
on  the  surface  and  their  passage  into  the  liquid. 
By  diffusion,  the  free  molecules  will  gradually 
move  away  from  the  neighborhood  of  the  surface 
of  the  solid  and  make  room  for  others,  and  thus, 
if  the  system  remains  undisturbed,  the  liquid  will  eventually  become 
a  solution  of  uniform  concentration.  If  a  large  enough  amount  of 
the  solid  has  been  provided,  the  ultimate  condition  will  be  that  of  a 
saturated  solution  with  excess  of  the  solid  beneath.  If  we  had  proper 
means  of  measuring  it,  the  tendency  of  the  molecules  to  leave  the 


FIG.  77. 


186  INORGANIC  CHEMISTRY 

solid  in  the  presence  of  a  given  liquid  would  give  the  effect  of  a  kind 
of  pressure.     This  is  spoken  of  as  solution  pressure. 

Now  the  molecules,  after  having  entered  the  liquid,  move  in 
every  direction,  and  consequently  some  of  them  will  return  to  the 
solid  and  attach  themselves  to  it.  The  frequency  with  which  this 
will  occur  will  be  greater  as  the  crowding  of  particles  in  the  liquid 
increases,  so  that  a  stage  will  eventually  be  reached  at  which  the 
number  of  molecules  leaving  the  solid  will  be  no  greater  than  that 
landing  upon  it  in  a  given  time.  If  the  whole  of  the  liquid  has  mean- 
while become  equally  charged  with  dissolved  molecules,  there  will  be 
no  chance  that  the  field  of  liquid  immediately  round  the  solid  will  lose 
them  by  diffusion,  so  that  a  condition  of  balance  or  equilibrium  (p. 
169)  will  have  been  established:  NaCl  (solid)  <=>  NaCl  (dslvd).  The 
presence  of  the  particles  in  the  liquid  produces  what  we  have  called 
diffusion  pressure;  and  when  the  diffusion  pressure,  by  the  continual 
increase  in  the  number  of  dissolved  molecules,  becomes  equal  to  the 
solution  pressure,  increase  in  concentration  of  the  solution  ceases.  It 
is  at  this  point  that  we  speak  of  the  solution  as  being  saturated  with 
respect  to  the  particular  substance  dissolving.  The  analogy  to  vapor 
tension  and  vapor  pressure  (p.  168)  is  evident. 

The  foregoing  explanation  should  be  compared  carefully  with  that 
given  in  the  section  on  the  molecular  relations  in  liquids,  and  in  that 
on  equilibrium  (pp.  168-170). 

%i  >f 

Solution  Pressure  and  Solution  Suction.  —  It  should  be 
noted  that  the  term  solution  pressure  really  presents  a  one-sided  view 
of  the  situation.  When  we  place  sulphur  in  carbon  disulphide,  it 
dissolves  freely,  and  we  should  therefore  say  that  sulphur  has  a  high 
solution  pressure.  But  when  we  put  sulphur  in  water,  hardly  any 
dissolves,  and  we  should  then  say  that  sulphur  had  almost  no  solution 
pressure.  Clearly  the  tendency  to  dissolve  is  a  mutual  property  of  the 
solvent  and  solute  together,  and  the  term  solution  pressure  is  misleading 
because,  when  taken  literally,  it  ascribes  the  result  entirely  to  the 
solute.  The  ability  of  the  solvent  to  take  up  the  solute,  which  we 
might  name  the  solution  suction  of  the  solvent,  is  at  least  equally 
important.  It  is  not  simply  a  question  of  how  strongly  the  sulphur 
ca^  throw  off  molecules,  but  also  of  how  liberally  the  solvent  is  able 
to  mix  with  these  molecules.  Water  is  not  able  to  take  up  or  to  hold 
much  sulphur,  but  carbon  disulphide  will  extract  or  cohere  with  a 
large  proportion  of  the  very  same  sulphur.  It  is  largely  a  question  of 
some  sort  of  physical  affinity  of  the  solvent  for  the  solute.  The  term 


SOLUTION  187 

solution  pressure  is,  of  course,  understood  by  physical  chemists  to 
include  this  affinity  or  miscibility,  but  the  words  are  far  from  convey- 
ing the  idea  to  the  student,  and  indeed  seem  explicitly  to  exclude  it. 
Unfortunately,  we  have  not  yet  been  able  to  discover  or  invent  any 
explanation  of  why  some  pairs  of  substances  are  so  freely  miscible, 
while  other  pairs  are  practically  immiscible. 

Independent  Solubility.  —  Just  as  two  gases,  when  mixed,  are 
independent  of  one  another  (p.  Ill),  and  have  severally  the  same 
pressure,  solubility,  and  so  forth,  as  they  would  possess  if  each  alone 
occupied  the  same  space,  so  it  is  with  dissolved  substances.  In  gen- 
eral, a  volume  of  water,  in  which  a  small  amount  of  some  substance 
has  been  dissolved,  will  take  up  as  much  of  a  second  substance  as 
would  an  equal  volume  of  pure  water.  Thus,  water  containing  some 
sugar  will  dissolve  as  much  sodium  chloride  as  the  same  amount  of 
pure  water.  In  the  kinetic-molecular  point  of  view,  the  dissolved 
molecules  of  sugar  have  no  connection  with,  or  influence  upon,  the 
mechanism  which  determines  the  solubility  of  the  salt,  namely,  the 
exchange  of  salt  molecules  between  the  suspended,  dissolving  crystals 
and  the  solution. 

Naturally  this  principle  of  independent  solubility  does  not  hold 
with  any  degree  of  exactness  when  the  concentration  of  the  substance 
already  present  is  great.  If  much  sugar  is  present,  then  the  solubility 
of  salt  in  the  mixture  will  depend  on  the  miscibility  of  salt  with  sugar, 
as  well  as  on  its  miscibility  with  water,  and  the  total  solubility  may 
be  very  different  from  that  in  water.  It  fails  also  when  the  two 
solutes  interact  chemically,  as  will  usually  be  the  case,  for  example, 
when  each  is  an  acid,  base,  or  salt  (see  Chap.  XIX).  The  solubility  is 
affected  in  an  especial  degree  when  the  two  substances  have  one 
radical  in  common,  as  when  they  are  nitric  acid  and  a  nitrate,  or  two 
chlorides  (see  Ionic  equilibrium). 

Conditions  Affecting  the  Solubility  of  a  Gas.  —  The  same 
conceptions  may  be  used  to  explain  any  case  of  solution.  Let  us  take 
that  of  oxygen  conducted  into  a  bottle  which  is  partially  filled  with 
water  (Fig.  78),  no  other  gas  being  present  in  the  space  above  the 
liquid.  As  the  molecules  of  the  gas  impinge  upon  the  liquid,  some 
of  them  pass  into  it  and  dissolve.  The  particles  which  have  thus 
gained  access  to  the  liquid  move  about  in  every  direction  and,  as  they 
become  more  and  more  numerous,  a  larger  and  larger  number  will 
escape  from  the  surface  and  pass  back  into  the  gaseous  condition. 


188  INORGANIC   CHEMISTRY 

At  first,  this  reaction  will  be  slight,  but  eventually,  as  the  solution 
increases  in  concentration,  it  must  become  equal  in  rate  to  the  process 
of  solution  itself.  It  is  assumed  that  the  supply  of  gas  is 
maintained  at  a  uniform  pressure,  and  therefore  uniform 
molecular  concentration,  during  the  whole  process.  Once 
more  we  shall  have  a  state  of  balance  or  equilibrium,  and 
the  liquid  will  be  saturated,  this  time  with  a  gas:  Oxygen 
(gas)  <=±  Oxygen  (dslvd) .  It  is  found,  as  the  hypothesis 
would  lead  us  to  expect,  that  the  concentration  of  the 
saturated  solution  of  any  given  gas  is  proportional  to  the 


pressure  at  which  the  gas  is  supplied,  that  is,  to  the  molecular 
concentration  of  the  gas  (Henry's  law). 

The  solubilities  of  different  gases,  even  when  their  pressures  are 
equal,  varies  much.  This  is  due,  of  course,  to  their  differing  physical 
affinities  for  (miscibilities  with)  the  liquid.  One  volume  of  water  will 
dissolve  1300  volumes  of  ammonia  at  0°  and  760  mm.,  while  it  will 
dissolve  only  about  0.02  volumes  of  hydrogen  under  the  same  condi- 
tions. In  one  volume  of  alcohol,  at  0°  and  760  mm.,  17.9  volumes  of 
hydrogen  sulphide  or  0.07  volumes  of  hydrogen  may  be  dissolved. 
Henry's  law  describes  the  behavior  of  any  one  gas  with  exactness  only 
when  low  gaseous  pressures,  or  gases  whose  solubility  is  small,  are  in 
question.  Great  solubility  must  be  due  in  part,  probably,  to  chemi- 
cal union  between  a  part,  at  least,  of  the  gas  and  the  solvent,  or  to 
cohesive  influences  which  the  molecules  of  the  dissolved  gas  exert 
upon  each  other.  The  law,  and  the  kinetic-molecular  explana- 
tion of  the  law,  on  the  other  hand,  consider  only  an  ideal  behavior 
involving  complete  chemical  and  physical  independence  of  the 
molecules. 

When  a  mixture  of  two  gases  is  in  contact  with  the  solvent,  the 
molecular  viewpoint  enables  us  successfully  to  foretell  what  will 
happen.  The  quantity  of  each  gas  which  can  remain  dissolved  must 
depend  simply  upon  the  frequency  with  which  its  own  molecules  strike 
the  liquid,  and  must  be  independent  of  the  presence  of  the  other  gas. 
Hence  the  solubility  of  each  gas  is  the  same  as  if  it  were  present  alone 
at  its  own  partial  pressure  (Dalton's  law,  p.  111).  In  other  words, 
each  gas  has  the  same  pressure,  and  therefore  the  same  solubility, 
as  it  would  possess  if  it  alone  occupied  the  whole  space  above  the 
liquid. 

Air  dissolving  in  water  is  an  illustration  of  this  principle.  It  does 
not  dissolve  as  a  whole,  but  the  oxygen  and  nitrogen  dissolve  each  in 
proportion  to  its  intrinsic  solubility  and  partial  pressure. 


SOLUTION  189 

It  is  easy,  by  the  use  of  this  law,  to  form  an  approximate  estimate  of  the  pro- 
portion of  oxygen  to  nitrogen  in  the  dissolved  gases.  The  air  may  be  taken  to  be 
at  760  mm.,  and  its  composition  by  volume  roughly  1/5  oxygen  and  4/5  nitrogen. 
The  separate  solubilities  of  the  gases  at  760  mm.  are,  respectively,  4  and  2  vol- 
umes in  100  volumes  of  water.  Their  partial  pressure  being  1/5  and  4/5  of  an 
atmosphere,  the  amounts  actually  dissolved  will  be  4  X  1/5  =  0.8  and  2  X  4/5  = 
1.6  in  100  volumes  of  water.  The  ratio  of  free  oxygen  to  nitrogen  in  the  water 
will  therefore  be  1:2.  Thus,  if  we  expel  the  dissolved  gases  by  boiling  the 
water,  we  find  that  they  contain  33  per  cent  oxygen  by  volume  while  air  contains 
only  21  per  cent. 

This  equilibrium,  Gas  (gaseous)  <=±  Gas  (dslvd),  can  be  reached, 
naturally,  from  the  other  direction,  namely  by  starting  with  a  solu- 
tion of  the  gas  and  a  space  above  the  solution  containing,  at  first, 
none  of  the  gas.  The  gas  leaves  the  solution  until  the  rates  of 
emission  and  return  become  equal.  Hence,  a  gas  may  be  entirely 
removed  from  solution  by  bubbling  a  foreign  gas  through  the  liquid. 
The  bubbles  furnish  the  space  to  receive  the  emitted  gas,  and  have 
a  large  surface,  so  that  the  process  goes  on  rapidly.  The  bubbles 
also  escape,  and  carry  with  them  the  emitted  gas,  so  that,  in  this 
case,  there  is  no  re-solution.  This  is  a  case  of  nullifying  one  of  the 
two  opposed  tendencies  (p.  170). 

Two  Immiscible  Solvents:  Law  of  Partition.  —  An  inter- 
esting application  of  the  same  ideas  may  be  made  to  a  case  which 
occurs  very  commonly  in  chemical  work.  If  we  shake  up  a  small 
particle  of  iodine  with  water,  we  find  that  it  dissolves  slowly,  giving 
eventually  a  saturated  but  very  dilute  solution.  If  now  ether  in 
sufficient  quantity  be  shaken  with  the  aqueous  solution,  the  greater 
part  of  the  iodine  will  find  its  way  into  the  ether,  and  be  contained 
in  the  brown  layer  which  rises  to  the  top.  The  process  of  removing 
a  substance  partially  from  solution  in  one  solvent  and  securing  it  in 
another  is  called  extraction.  We  find  in  such  cases  that  neither 
solvent  can  entirely  deprive  the  other  of  the  whole  of  the  dissolved 
substance,  if  the  latter  is  soluble  in  both  independently:  I  (in  Aq)  ^±  I 
(in  ether).  The  partition  of  the  substance  takes  place  in  proportion 
to  its  solubility  in  each  solvent.  It  is  found  that  any  amount  of  the 
solute,  up  to  the  maximum  the  system  can  contain,  provided  this 
does  not  involve  too  high  a  concentration  in  either  solvent,  is  divided 
so  that  the  ratio  of  the  concentrations  in  the  two  solvents  is  always 
the  same  (Law  of  partition).  In  the  case  of  iodine  divided  between 
water  and  ether,  this  ratio  is  about  1  :  200. 


190  INORGANIC  CHEMISTRY 

The  aqueous  solution  of  potassium  iodide  has  a  very  great  power  of  dis- 
solving iodine,  and  we  find  that  in  the  presence  of  this  salt  the  ether  leaves  a  much 
larger  share  of  the  element  in  the  lower  layer.  A  part  of  the  iodine  combines  to 
form  KI3,  however,  so  that  this  is  not  a  case  of  simple  solution,  and  the  law  of 
partition  does  not  hold.  The  chemical  equilibrium  (see  Chap.  XV)  between  free 
and  combined  iodine  in  the  aqueous  layer  has  to  be  considered. 

This  principle  is  used  in  Parke's  process  (q.v.)  for  extracting 
silver  from  molten  lead,  by  means  of  melted  zinc  as  the  second 
solvent.  It  is  employed  in  separating  interesting  compounds  from 
animal  secretions  and  vegetable  extracts,  and  in  purifying  such 
compounds.  Nicotine  from  tobacco  and  cocaine  from  coca-leaves, 
are  secured  in  this  way. 

Influence  of  Temperature  on  Solubility.  —  The  quantity  of 
a  substance  which  we  can  dissolve  in  a  fixed  amount  of  a  given  solvent 
depends  very  largely  upon  the  temperature  of  both.  Usually  the 
solubility  increases  with  rise  in  temperature.  Measurements  may  be 
made  by  the  method  described  before  (p.  180),  using  excess  of  the 
finely  powdered  solute  with  different  portions  of  the  same  solvent  in 
vessels  kept  at  different  temperatures.  The  most  useful  way  of  repre- 
senting the  results  is  to  plot  them  graphically.  The  diagram  (Fig.  79) 
shows  the  curves  for  a  few  familiar  substances.  The  ordinates  repre- 
sent the  number  of  grams  of  the  anhydrous  compound  which  is  held 
in  solution  by  100  g.  of  water  in  each  case.  The  abscissae  represent 
the  temperatures.  The  concentration  for  any  temperature  can  be 
read  off  at  once.  Thus  100  g.  of  water  holds  13  g.  of  potassium  nitrate 
in  solution  at  0°  and  150  g.  at  73°.  The  increase  in  solubility  is  here 
enormous.  On  the  other  hand,  the  same  quantity  of  water  will  hold 
35.6  g.  of  sodium  chloride  in  solution  at  0°  and  39  g.  at  100°.  The 
difference  is  shown  at  once  when  we  examine  the  curves  and  observe 
that  the  line  representing  the  solubility  of  sodium  chloride  scarcely 
rises  at  all  between  0°  and  100°,  while  that  of  potassium  nitrate  is 
extremely  steep. 

Cases  in  which  the  solubility  decreases  with  rise  in  temperature  are 
less  common.  The  solubility  of  slaked  lime  (calcium  hydroxide 
Ca(OH)2,  used  to  make  limewater)  is  0.175  g.  at  20°  and  0.0789,  or 
less  than  half  as  great,  at  100°.  Anhydrous  sodium  sulphate  Na2SO4 
(Fig.  80,  p.  193)  is  another  illustration.  When  triethylamine,  an 
organic  base,  N(C2H5)3,  liquid  at  ordinary  temperatures,  is  added  to 
cold  water  until  no  more  will  dissolve,  the  solution,  which  is  perfectly 
clear  and  transparent,  on  being  warmed  with  the  hand  at  once  be- 


SOLUTION 


191 


192  INORGANIC  CHEMISTRY 

comes  clouded  from  the  separation  of  the  two  liquids.  A  compara- 
tively slight  elevation  in  temperature  causes  a  separation  into  two 
distinct  layers. 

Phases.  —  We  can  frequently  abbreviate  our  statements  by  using 
two  words  of  broad  significance,  one  of  which  has  already  been  em- 
ployed. A  set  of  materials  in,  or  tending  towards,  a  condition  of  equi- 
librium is  called  a  system.  The  discrete  parts  of  an  inhomogeneous 
system  are  called  its  phases.  Thus,  a  liquid  with  its  vapor  forms  a 
system  with  a  liquid  phase  and  a  vapor  phase.  A  saturated  solution 
is  a  system  with  three  phases,  the  undissolved  excess  of  the  solute 
(solid  phase),  the  solution,  and  the  vapor. 

Equilibrium  in  a  Saturated  Solution .  —  Once  a  solution  has 
become  saturated,  the  undissolved  material  remains  thereafter  un- 
changed in  amount,  no  matter  how  long  the  materials  are  left  in 
contact.  In  technical  terms,  the  quantity  of  each  phase  has  no  in- 
fluence on  the  concentration  of  any  of  them.  A  greater  excess  of  the 
solute  forces  no  more  matter  into  solution  than  does  a  small  excess. 

It  should  be  clearly  understood,  however,  that  the  kinetic  theory 
requires  us  to  assume  that  an  exchange  of  molecules  (p.  186)  is  still 
going  on  between  the  solid  and  the  solution.  That  this  conception  is 
correct  may  be  shown  in  various  ways.  Thus,  if  a  crystal,  the  edges 
or  corners  of  which  have  been  broken,  is  suspended  in  a  saturated 
solution  of  the  same  substance,  it  neither  increases  nor  diminishes  in 
weight.  Yet  we  find  that  the  imperfections  are  removed,  and  that 
this  takes  place  by  the  solution  of  a  portion  of  the  substance  from  the 
perfect  surfaces  and  its  deposition  upon  the  imperfect  ones. 

Supersaturated  Solutions.  —  Another  very  striking  proof  of 
this  may  be  obtained  by  saturating  water  with  ordinary  Glauber's 
salt  (hydrated  sodium  sulphate,  Na^SO^lOH^O)  at  a  temperature 
somewhat  above  the  ordinary,  say  30°,  at  which  temperature  100  c.c. 
of  water  hold  in  solution  40  g.  of  Na^SC^  (Fig.  80).  The  excess  of  the 
solid  is  carefully  and  completely  separated  from  the  liquid,  and  the 
latter  is  allowed  to  cool,  say  to  15°  in  a  flask  loosely  stoppered  with 
cotton.  The  solution  now  contains  (Fig.  80)  a  much  larger  amount  of 
sodium  sulphate  (Na^SO^  than  at  its  present  temperature  it  could 
acquire  from  contact  with  Glauber's  salt  (13  g.  at  15°).  Yet  in  the 
absence  of  a  crystal,  with  which  the  above  described  exchange  could 
take  place,  no  deposition  of  the  dissolved  substance  begins.  The 


SOLUTION 


193 


solution  may  be  kept  indefinitely  without  alteration.  The  intro- 
duction, however,  of  the  minutest  fragment  of  the  decahydrate  at 
once  starts  the  exchange,  and  this  is  necessarily  very  much  to  the 


0°   10°  20°   30°  40°   50°   60°   70°  80°  90°  100» 
Temperature 

FIG.  80. 


disadvantage  of  the  solution  and  the  advantage  of  the  crystal: 
Na2S04  (dslvd)  +  10H2O  <±  Na2SO4,10H20  (solid).  The  latter 
therefore  forms  the  center  of  a  radiating  mass  of  blade-like 
processes,  which  sprout  with  astonishing  rapidity 
through  the  liquid  (Fig.  81). 

Usually  the  cooling  of  a  concentrated  solution 
leads  to  the  almost  immediate  appearance  of  crystals 
spontaneously,  and  the  substance  is  deposited  gradu- 
ally as  the  temperature  falls.  But  solutions  of  a 
number  of  common  substances,  such  as  sodium 
thiosulphate  (photographer's  "hypo")  and  sodium 
chlorate,  behave  like  that  of  sodium  sulphate.  They 
are  said  to  have  a  tendency  to  give  supersaturated 
solutions.  In  general,  crystallization  can  be  started 
only  by  introduction  of  a  specimen  of  the  same 
substance,  or  at  all  events  of  one  isomorphous  (q.v.) 
with  it.  The  smallest  particle  of  the  right  material 
floating  in  the  air,  if  it  gains  accidental  admission,  will  bring  about  the 
result.  This  shows  the  importance  of  the  interchange  of  molecules 
of  which  we  have  spoken  for  establishing  equilibrium. 


Fia.  81. 


194  INORGANIC  CHEMISTRY 

Metastable  Condition.  —  The  above  phenomenon  is  not  an 
isolated  or  exceptional  one  in  physical  science.  It  is  commonly  the 
case  that,  when  the  conditions  for  some  physical  change  have  been 
reached,  the  beginning  of  the  physical  change  is  delayed  or  entirely 
fails.  The  system  is  then  said  to  be  in  a  metastable  condition.  Un- 
stable it  is  not.  Yet  it  is  not  in  the  state  of  greatest  stability,  for  the 
element  of  equilibrium  is  lacking.  Thus,  pure  water  may  easily  be 
cooled  three  or  four  degrees  below  0°  without  the  appearance  of  any 
ice.  Agitation,  however,  in  this  case,  results  in  the  appearance  of  ice 
sooner  or  later.  Addition  of  a  fragment  of  ice  results  in  immediate 
crystallization  (freezing).  In  like  manner  water  may  be  heated  to  a 
temperature  above  100°  without  boiling.  Drops  of  water,  suspended 
in  some  oil  of  almost  the  same  specific  gravity,  may  even  be  raised  to 
175°  before  the  water  turns  into  steam.  This  is  because  there  is  no 
water  vapor  (but  only  oil)  in  contact  with  the  water.  Similarly,  air 
which  is  saturated  with  moisture,  if  it  contains  no  dust,  may  be 
cooled  without  the  appearance  of  fog. 

Phenomena  of  supersaturation,  of  a  temporary  kind  at  least,  are  extremely 
common  in  chemistry.  Almost  every  delayed  precipitation  is  a  case  of  it.  Ba- 
rium sulphate,  for  example,  is  always  slow  in  appearing  in  dilute  solutions.  So  is 
sulphur,  set  free  from  dilute  sodium  thiosulphate  solution  by  the  action  of  an  acid. 
In  the  latter  case,  instant  reneutralization  with  a  base  does  not  prevent  the 
ultimate  appearance  of  the  sulphur,  showing  that  the  cause  does  not  lie  in  slow 
interaction  of  the  salt  with  the  acid. 

Definition  of  a  Saturated  Solution:  A  Warning.  —  To  avoid 
a  common  misconception,  it  must  be  noted  that  a  saturated  solution 
must  not  be  defined  as  one  containing  all  of  the  solute  that  it  can 
hold.  A  supersaturated  solution  holds  more.  The  saturated  solution 
is  one  which  contains  all  of  the  dissolved  solute  that  it  can  acquire 
from  the  undissolved  solute.  Better  still,  it  is  that  solution  which, 
when  placed  with  excess  of  the  solute,  is  found  to  be  in  equilibrium. 

It  must  be  clearly  understood  that  solution  is  not  a  process  of 
filling  the  pores  of  the  liquid.  If  that  were  true,  approximately 
equal  weights  of  all  substances  would  find  accommodation  in  equal 
volumes  of  water.  The  fact  is  that,  for  example,  100  c.c.  of  water 
can  dissolve  195  g.  of  silver  fluoride,  but  only  0.00000035  g.  of  silver 
iodide,  although  the  space  available  in  the  solvent  (if  there  is  any 
free  space)  is  the  same  in  both  cases. 

The  same  conclusion  is  reached  when  we  consider  that  two  forms 
of  the  same  compound  have  different  solubilities.  Thus,  at  20°, 


SOLUTION  195 

Na2S04,10H20  can  give  about  18  g.  of  Na^SCU  to  100  c.c.  of  water 
(Fig.  80).  But  anhydrous  sodium  sulphate  Na2SO4  at  20°  gives 
59  g.  to  the  same  amount  of  water  (read  up  to  dotted  line,  Fig.  80). 

Saturation.  —  When  we  have  shaken  a  solid  for  a  sufficient 
length  of  time  with  a  given  amount  of  a  liquid,  we  obtain  a  solution 
which  is  saturated  with  respect  to  that  substance.  Having  called 
this  a  saturated  solution,  we  are  inclined  to  extract  from  the  term  a 
meaning  different  from  that  which  it  was  really  intended  to  convey. 
We  are  in  danger  of  thinking  that  the  solution  itself  is  in  some  way 
peculiar  —  that,  for  instance,  it  contains  all  of  the  solid  which  it  is 
capable  of  holding.  This  would  be  an  entire  misconception.  If  we 
desire  to  make  a  solution  of  sodium  sulphate  Na^SO^  for  example,  we 
may  present  this  substance  to  the  water  either  in  the  form  of  Glau- 
ber's salt  Na2SO4,10H20  or  of  anhydrous  sodium  sulphate.  Now  the 
anhydrous  and  the  hydrated  forms  of  a  substance  always  behave  like 
entirely  different  substances.  This  hydrate  cannot  give  more  than  5 
parts  of  sodium  sulphate  Na^SC^  to  100  parts  of  water  at  0°.  When 
the  anhydrous  compound  is  used,  many  times  this  amount  (Fig.  80)  is 
dissolved  at  the  same  temperature :  Na2S04  (solid)  +±  Na2SO4  (dslvd) . 
The  solution  pressures  of  the  two  forms  are  entirely  different.  The 
phrase,  "a  saturated  solution  of  sodium  sulphate,"  is  therefore  devoid 
of  definite  meaning.  We  must  describe  the  liquid  as  a  solution  satu- 
rated by  anhydrous,  or  by  hydrated  sodium  sulphate  as  the  case 
may  be. 

Being  different  substances,  the  hydrated  and  anhydrous  forms  of  a  compound 
must  be  investigated  separately  as  to  their  solubility  at  various  temperatures. 
The  results  must  give  different  curves,  as  for  distinct  substances. 

Before  referring  to  the  curves  of  the  two  sodium  sulphates,  it  must  be  re- 
marked that  hydrates  decompose  into  the  anhydrous,  or  some  less  hydrated  form 
at  a  definite  temperature.  We  cannot  therefore  continue  the  observation  of  the 
solubility  of  the  substance  beyond  the  temperature  at  which  it  ceases  to  exist. 
Thus  the  solubility  curve  of  a  hydrate  comes  to  an  abrupt  termination  at  the 
decomposition,  or,  as  it  is  usually  called,  the  transition  point.*  Now  the  decahy- 
drate  of  sodium  sulphate  melts  and  decomposes  at  32.4°,  so  that  its  solubilities  can 
be  measured  only  from  about  0°  to  32.4°.  The  solubility  of  the  anhydrous  form, 
however,  can  be  investigated  up  to  100°,  or  beyond  it,  if  necessary.  Not  only  so, 
but  measurements  can  be  carried  cut  below  32.4°.  There  is  time  to  saturate  the 

*  These  statements  refer  to  equilibrium  conditions.  In  practice,  by  rapid 
work,  measurement  of  the  solubilities  of  hydrates  can  usually  be  made  above 
the  melting  (transition)  point  also. 


196  INORGANIC  CHEMISTRY 

solution  with  the  anhydrous  body,  and  decant  the  liquid  for  analysis,  before  the 
hydrate  begins  to  crystallize  out. 

The  solubility  of  the  decahydrate  (Fig.  80)  rises  rapidly  between  0°  and  32.4° 
from  5  to  55  parts  in  100.  The  solubility  of  the  anhydrous  sodium  sulphate  de- 
creases steadily  from  more  than  55  parts  below  32.4°  to  42.5  parts  at  100°.  The 
character  of  the  two  bodies,  in  the  matter  of  solubility,  is  therefore  entirely  differ- 
ent. The  solutions  themselves,  as  has  been. said  already,  are  identical  in  every 
way,  when  they  have  the  same  concentration,  whether  they  have  been  made  from 
the  one  substance  or  the  other. 

Let  us  return  now  to  the  proper  use  of  the  term  "saturated  solution."  We 
might  say,  correctly,  that  at  20°  a  solution  made  from  hydrated  sodium  sulphate 
and  containing  19.4  parts  of  sodium  sulphate  in  100  of  water  was  saturated.  This 
would  not,  however,  be  the  maximum  quantity  which  the  same  amount  of  water 
could  hold,  for,  with  the  help  of  the  anhydrous  compound,  we  could  add  an 
amount  equivalent  to  prolonging  the  ordinate  at  20°  until  it  intersected  the  curve 
of  anhydrous  sodium  sulphate  somewhere  about  the  value  60.  Nor  "can  we  be 
sure  that  even  then  the  water  would  contain  all  the  sodium  sulphate  which  it 
could  hold.  It  is  conceivable  that,  by  presenting  the  substance  in  some  still  other 
form,  even  greater  solubility  might  be  observed. 

A  saturated  solution,  if  we  fix  our  minds  upon  it  simply  as  a  solution,  is  not 
different  from  any  other  solution.  There  is  no  feature  in  the  properties  of  such 
a  solution,  qua  solution,  which  distinguishes  it  in  the  least  from  one  containing 
slightly  less  or  one  containing  slightly  more  of  the  dissolved  substance.  In  con- 
tact with  a  crystal  of  the  substance  used  to  produce  the  saturation,  however,  the 
saturated  solution  is  found  to  be  in  equilibrium,  while  the  unsaturated  solution 
takes  up  more  of  the  substance,  and  the  supersaturated  solution  at  once  deposits 
the  amount  which  it  contains  in  excess  of  the  saturated  solution.  The  words  un- 
saturated, saturated,  and  supersaturated  convey,  therefore,  no  meaning  unless  we 
add  that  the  solution  is  so  towards  some  specific  form  of  material.  These  are 
qualities  of  the  system  including  the  undissolved  body,  and  not  of  the  solution  by 
itself. 

The  fact  that  the  hydrated  and  anhydrous  sodium  sulphates  give  saturated 
solutions  of  the  same  concentration  at  32.4°,  so  that  the  curves  intersect  at  this 
point,  is  a  very  significant  one.  This  is  the  temperature  at  which  the  former  sub- 
stance turns  into  the  latter.  At  transition  points  like  this  the  values  of  solubility, 
vapor  pressure,  and  some  other  properties,  are  always  the  same  for  both  forms 
(see  Freezing-points  of  solutions). 

The  Solute  Modifies  the  Physical  Properties  of  the  Sol- 
vent. —  These  alterations  in  the  values  of  the  physical  properties 
of  the  solvent  are  divisible  into  two  classes.  In  one  of  these  classes, 
equal  numbers  of  dissolved  molecules  of  different  substances,  dissolved 
in  equal  quantities  of  the  solvent,  produce  the  same  amount  of  change 
in  the  properties  affected.  The  change  is  proportional  to  the  molecu- 
lar concentration  of  the  dissolved  body,  when  this  concentration  is 


SOLUTION  197 

small.  This  effect  appears,  therefore,  to  be  largely  due  to  mechan- 
ical causes,  since  it  depends  simply  upon  the  number  of  foreign 
molecules,  without  reference  to  the  substance  composing  them.  Of 
this  nature  are  the  lowering  in  the  vapor  tension  of  the  liquid,  the 
raising  of  its  boiling-point,  the  lowering  of  the  freezing-point  of  the 
liquid  (see  below),  and  the  value  of  the  osmotic  pressure  (see 
Chap.  XVII). 

In  the  other  class,  the  extent  of  the  effect  varies  in  amount  with 
the  nature  of  the  substance  dissolved,  even  when  equal  numbers  of 
molecules  are  used.  This  sort  of  modification  in  the  properties  is 
specific  for  each  pair  of  substances  (solvent  and  solute).  The  change 
in  total  volume  occurring  during  solution,  and  the  heat  of  solution 
(see  below)  belong  to  this  class. 

The  Vapor  Tensions  of  Solutions.  —  A  solute,  which  is  itself 
non-volatile,  tends  to  diminish  the  vapor  tension  of  the  solution.  It 
hinders  the  emission  of  the  vapor.  Thus  if,  instead  of  water,  we 
introduce  aqueous  solutions  of  differing  concentrations  of  the  same 
substance  successively  into  the  barometric  vacuum  (Fig.  53,  p.  145), 
we  find  that  the  vapor  pressures  of  the  solutions  are  less  than  that  of 
water  at  the  same  temperature.  The  diminution  in  the  fall  of  the 
mercury  column,  which  measures  the  lowering  in  vapor  pressure,  is 
proportional  to  the  concentration  of  each  solution.  The  limit  is 
reached  with  the  saturated  solution,  although,  if  this  is  rather  con- 
centrated, the  proportionality  does  not  hold  strictly  down  to  that 
point  (see  Chap.  XVII).  For  example,  at  20°,  water  has  the  vapor 
pressure  17.4  mm.  while  a  saturated  solution  of  common  salt  (34.3  : 
100  Aq)  shows  only  13.9  mm.  —  a  lowering  in  the  vapor  pressure  of 
the  solvent  by  3.5  mm. 

If  the  substance  is  very  soluble,  and  the  solution  highly  concen- 
trated, the  lowering  in  the  vapor  tension  will  be  considerable.  In 
fact,  the  solution  may  give  a  vapor  pressure  of  water  less  than  that 
commonly  present  in  the  atmosphere.  Such  a  solution,  placed  in  an 
open  vessel,  will  not  evaporate.  On  the  contrary,  vapor  from  the  air 
will  enter  it,  and  it  will  increase  in  bulk.  For  this  reason,  crystals  of 
very  soluble  substances  are  usually  moist  and,  when  exposed  to  the 
air,  acquire  water  from  the  latter  and  dissolve  in  this  water.  This 
behavior  is  called  deliquescence,  and  is  exhibited,  for  example,  by  the 
hydrate  of  calcium  chloride  CaCl2,2H2O,  which  is  consequently  used 
for  drying  gases.  Magnesium  chloride  MgC^,  present  as  an  impurity 
in  common  salt,  causes  the  latter  to  become  moist  in  damp  weather. 


198 


INORGANIC  CHEMISTRY 


FIG.  82. 


The  principle  involved  will  become  clear  if  we  imagine  two  vessels, 
one  containing  pure  water  and  one  an  aqueous  solution,  to  be  placed 

on  a  glass  plate  and  covered  by  a  bell  jar 
(Fig.  82).  Each  liquid  exchanges  water 
molecules  with  the  moist  air  in  the  jar, 
but  the  solution  gives  off  water  more  feebly 
than  does  the  pure  water.  The  result  is 
that  the  latter  can  produce  a  pressure  of 
water  vapor  higher  than  that  which  would 
be  in  equilibrium  with  the  solution.  The 
solution,  therefore,  receives  continuously 
more  molecules  than  it  emits,  and  increases  in  volume.  The  pure 
water  thus  gradually  passes  through  the  vapor  state  into  the  solu- 
tion until  it  is  all  gone.  If  sufficient  water  was  present,  the  process 
would  go  on  until  the  solution  became  infinitely  dilute. 

One  of  the  oddities  of  chemical  literature  is  the  fact  that  most  authors  dis- 
cuss efflorescence  (p.  151)  and  deliquescence  in  the  same  paragraph,  if  not  in 
the  same  sentence.  Since  these  two  topics  are  entirely  unrelated,  and  belong 
indeed  to  different  chapters,  this  inappropriate  juxtaposition  is  somewhat  con- 
fusing. 

Elevation  in  the  Boiling-Point  of  a  Solution.  —  A  solute 
lowers  the  vapor  pressure  of  a  solvent,  not  only  at  room  temperature, 
but  at  all  temperatures.  Thus,  at  100°,  a  7.5  per  cent  solution  of 
potassium  chloride  shows  a  mm 
vapor  pressure  of  only  734.1  7eo 
mm.,  while  that  of  water  is 
760  mm.  The  lowering  is 
25.9  mm.  Hence,  the  solu-  _. 
tion  has  to  be  raised  to  a  i 
higher  temperature  (100.96°)  | 
before  it  boils.  This  is  al-  o 
most  exactly  0.037°  per  1  | 
mm.,  which  is  the  value  for 
pure  water  (p.  145). 


Temperature 
FIG.  83. 


This  conclusion  may  also  be 
reached  graphically  (Fig.  83). 
The  ordinates  represent  the 
vapor  tensions  corresponding  to  the  temperatures  shown  by  the  abscissae. 
They  increase  in  length  with  rise  in  temperature.  The  horizontal  dotted 
line  shows  the  vapor  tension  of  760  mm.  at  which  any  liquid  will  boil.  The 


SOLUTION  199 

boiling-point  of  the  solvent  is  therefore  the  temperature  at  which  its  curve 
intersects  this  line.  Since  the  vapor  tensions  of  the  solution  are  all  below  those 
of  the  solvent,  its  curve  lies  below  that  for  the  solvent,  but  ascends  along  with 
the  latter.  Hence,  it  also  cuts  the  760  mm.  line,  but  at  a  point  beyond  the  boil- 
ing-point of  the  pure  solvent.  Since,  for  short  lengths,  the  curves  are  very  nearly 
straight  lines,  the  distances  from  boiling-point  to  boiling-point  are  very  nearly 
proportional  to  the  vertical  distances  between  the  curves.  That  is  to  say,  the 
elevations  in  the  boiling-point  are  proportional  to  the  depressions  in  the 
vapor  tension,  and  therefore  to  the  concentration  of  the  different  solutions 
of  any  one  substance". 

When  different  substances  in  the  same  solvent  are  compared,  equal  mo- 
lecular concentrations  of  the  solute  give  equal  elevations  in  the  boiling- 
point.  Thus,  1  gram-molecular  weight  of  sugar  (342  g.),  or  of  glycerine  (92  g.), 
dissolved  in  1000  cc.  of  water,  will  elevate  the  boiling-point  from  100°  to  100.52° 
(for  exceptions,  see  Chap.  XVII).  When  the  effects  of  the  same  solute  in  differ- 
ent solvents  are  compared,  the  scale  of  the  elevation  alone  is  different. 

The  same  effect  may  be  observed  in  isomorphous  mixtures  of  solid  bodies, 
which  chemists  consider  to  be  solid  solutions.  Thus  the  vapor  tension  of  water 
in  the  alums  is  greater  than  the  average  vapor  pressure  of  water  in  the  air,  and 
hence  they  lose  their  water  of  hydration  spontaneously  (cf.  p.  152).  But  if  mixed 
crystals  of  two  alums,  say  ordinary  alum  K2SO4,A12(SO4)3,24H2O  and  iron  alum 
(NH4)2SO4,Fe2(SO4)3,24H2O  are  prepared,  they  keep  perfectly.  The  vapor  ten- 
sion of  the  water  in  each  has  been  lowered  by  the  influence  of  the  other  alum 
dissolved  in  it.  Similarly,  calcium  formate  Ca(CHO2)2,4H2O  effloresces  (p.  151), 
but  loses  this  tendency  when  crystallized  with  some  of  the  isomorphous  barium  or 
strontium  salt. 

Freezing-Points  of  Solutions:  Freezing  Mixtures.  —  Every 
pure  liquid  has  a  definite  temperature  at  which  it  freezes.  Thus, 
pure  water  freezes  at  0°  and  benzene  at  5.5°.  The  freezing  of  a  dilute 
solution  consists,  usually,  in  the  crystallization  of  some  of  the  pure 
solvent  only.  The  presence  of  a  foreign,  dissolved  body,  tends  to 
prevent  this  freezing,  and  so  solutions  can  be  frozen  only  at  tem- 
peratures below  the  freezing-point  of  the  pure  solvent.  Thus,  sea- 
water  is  harder  to  freeze  than  fresh  water.  Also,  the  ice  formed  in  salt 
water  is  free  from  salt.  The  freezing-points  of  solutions  of  the  same 
substance  are  found  to  be  depressed  below  that  of  the  solvent  in  pro- 
portion to  the  concentration  of  the  solute  *  (see  Chap.  XVII).  This 
may  be  shown  graphically  (Fig.  84).  The  ordinates  represent  vapor 
tensions  corresponding  to  the  temperatures  shown  by  the  abscissae. 
The  rate  at  which  the  former  rise  is  greater  for  ice  than  for  water, 

*  Only  when  the  solid  which  separates  consists  of  the  pure  solvent,  and  the 
solute  does  not  enter  into  it,  does  this  law  represent  the  facts. 


4.6 


200  INORGANIC  CHEMISTRY 

hence  the  ice  curve  is  steeper.  At  0°,  ice  and  water  can  coexist 
permanently  (p.  144).  By  measurement,  they  have  the  same  vapor 
tension  (4.6  mm.)  at  this  point.  Theoretically,  if  they  had  not,  they 
could  not  coexist  indefinitely,  for  the  one  with  the  greater  vapor 

tension  would  evaporate, 
and  its  vapor  would  con- 
dense on  the  other  until 
one  of  them  alone  re- 
mained. 

Now  the  vapor  ten- 
sion of  a  solution  is,  at 
all  temperatures,  lower 
(Fig.  84)  than  that  of 

„ , ,  water.    Hence,  fora  solu- 

Temp-  °°  tion,  the  curve  must  cut 

the  ice  curve  below  4.6 

mm.  and  therefore  behind  0°.  In  other  words,  ice  and  the  solution 
cannot  have  equal  vapor  tensions,  and  therefore  coexist  indefinitely, 
except  at  some  temperature  below  0°.  But  the  temperature  at 
which  ice  can  exist  indefinitely  in  a  solution  is  the  freezing-point. 
Hence,  freezing-points  of  solutions  are  always  lower  than  those  of 
the  pure  solvents. 

By  measuring  the  vapor  tensions  of  the  solution  at  several  temper- 
atures, in  order  to  see  how  far  the  curve  for  the  solution  is  below  that 
of  water,  and  then  producing  the  curve  for  the  solution  backwards,  the 
intersection  with  the  ice  curve,  and  therefore  the  freezing-point,  may 
be  obtained  graphically.  Direct  measurement  always  confirms  the 
result. 

Since  the  ice  curve  is,  for  a  short  distance,  almost  a  straight  line,  it 
follows  that  the  depressions  of  the  freezing-point  are  proportional  to 
those  of  the  vapor  tension,  and  these  in  turn  are  proportional  to  the 
concentrations.  From  this  relation  we  get  the  statement  with  which 
this  section  opened.  Thus,  solutions  of  sugar  containing  11.4,  22.8, 
and  34.2  g.  of  sugar  to  100  g.  of  water  freeze  at  -0.62°,  -1.23°,  and 
—  1.85°,  respectively.  Numerically,  in  the  case  of  water,  a  lowering 
of  the  vapor  pressure  by  jiir  of  its  amount  at  each  temperature  sets 
the  freezing-point  back  1.05°. 

This  fact  likewise  explains  why  salt,  thrown  on  ice,  causes  the 
latter  to  melt.  A  saturated  solution  of  salt  does  not  freeze  until 
cooled  to  —21°  (  —  6°  F.),  and  it  then  gives  a  mixture  of  pure  ice  and 
pure  salt  crystals,  known  as  a  cryohydrate.  Hence,  ice  and  salt  can 


SOLUTION  201 

not  permanently  exist  together  above  —  21°,  their  cryohydratic  point. 
Below  —  6°  F.,  salt  will  no  longer  melt  ice  because  the  solid  mixture 
can  exist  below  that  temperature.  A  mixture  of  ice  and  salt,  giving 
a  temperature  of  —  6°  F.,  is  called  a  freezing  mixture,  and  is  used  in 
freezing  ice  cream  and  ices.  Substances  more  soluble  than  common 
salt,  such  as  calcium  chloride,  give,  with  ice,  freezing  mixtures  of 
lower  temperature  (see  end  of  section  on  Heat  of  solution,  p.  203). 

Densities  of  Solutions.  —  The  density  of  a  solution  is  usually 
greater,  though  often  less,  than  that  of  water  and,  in  each  case,  varies 
with  the  concentration.  For  commercial  purposes,  the  concentra- 
tion of  a  solution  is  commonly  defined  by  the  density  (or  specific 
gravity).  Thus,  we  purchase  ammonium  hydroxide  solution  of 
"0.88  sp.  gr.,"  meaning  35  per  cent  of  ammonia,  or  sulphuric  acid  of 
"  1.84  sp.  gr.,"  meaning  94.8  per  cent  of  the  acid. 

The  commonly  greater  density  of  a  solution  is  utilized  in  making 
solutions  in  chemical  factories.  Shaking  several  tons  of  the  mixture 
is  out  of  the  question,  and  stirring  costs  money.  If  the  solid  is 
placed  in  the  bottom  of  the  tank,  under  water,  a  saturated  solution  is 
formed  in  the  lowest  layer  of  the  water,  and  passage  of  the  dissolving 
substance  into  the  upper  layers,  by  diffusion,  would  take  months  or 
years.  Hence  most  of  the  solid  would  remain  undissolved  (Fig.  76, 
p.  184).  But  when  the  solid  is  placed  on  a  shelf,  near  the  surface  of 
the  water,  the  solution,  being  more  dense,  sinks  through  the  water, 
fresh  water  rises  to  the  shelf,  and  a  circulation  is  started.  This  results 
in  the  dissolving  of  the  whole  material  in  a  surprisingly  short  time, 
with  no  expenditure  of  labor  whatever. 

Changes  in  Volume  upon  Solution.  —  The  erratic  and,  at 
present,  unexplained  changes  in  volume  which  occur  when  a  substance 
is  dissolved,  seem  to  indicate  that  the  process  is  less  simple  than  we 
have  thus  far  admitted,  and  that  chemical  changes  occur  during  the 
process.  Thus,  when  250  g.  of  common  salt  are  dissolved  in  1  liter 
of  water  (=  1000  c.c.  =  1000  g.),  which  gives  a  20  per  cent  solution, 
the  volume  of  the  solution  is  only  1086  c.c.  Since  the  250  g.  of  salt 
occupied  116  c.c.  before  being  dissolved,  a  shrinkage  of  1116-1086,  or 
30  c.c.,  accompanied  the  process  of  solution.  On  the  other  hand, 
214  g.  of  ammonium  chloride  (volume  142.5  c.c.)  and  843.5  c.c.  of 
water  have  a  total  volume  of  986  c.c.,  but  when  dissolved  give 
1000  c.c.  of  solution.  Here  there  is  an  expansion  of  14  c.c.  Table 
sugar,  however,  dissolves  in  water  with  almost  no  change  in  volume. 


202  INORGANIC  CHEMISTRY 

Is  Solution  a  Physical  or  a  Chemical  Change?  —  These 
phenomena  are,  in  part,  accounted  for  by  the  fact  that  water  is  not  a 
single  substance,  but  a  mixture.  It  is  largely  composed  of  dihydrol 
(H20)2,  with  much  trihydrol  (H20)3  near  to  0°  and  increasing  quanti- 
ties of  monohydrol  H2O  at  higher  temperatures.  When  any  sub- 
stance is  dissolved  in  considerable  amount  in  water,  the  equilibrium 
amongst  these  three  kinds  of  molecules  is  disturbed,  and  their  pro- 
portions change: 

2(H2O)8  ^  3(H20)2  ^  6H2O. 

Now,  equal  weights  of  these  three  kinds  of  water  occupy  different 
volumes,  and  hence  solution  is  accompanied  by  changes  in  the  volume 
of  the  water.  The  same  condition  in  water  explains  the  point  of 
maximum  density  (4°).  The  change  from  (H2O)3  to  (H2O)2,  which 
proceeds  as  the  temperature  rises  from  0°  to  4°,  is  accompanied  by  a 
shrinkage,  because  dihydrol  has  the  higher  specific  gravity.  Beyond 
4°,  the  usual  expansion  with  rising  temperature  prevails. 

There  is  also  evidence  that  many  dissolved  bodies  form  unstable 
compounds  with  water,  although  we  have  not  as  yet  definite  infor- 
mation about  these  compounds. 

Dissolving  in  water  is,  therefore,  partly  a  chemical  and  only 
partly  a  physical  process  —  a  part  of  the  water  is  always  affected,  and 
a  part  or  all  of  the  solute  may  go  into  combination. 

Heat  of  Solution.  —  As  in  other  changes  of  state  of  aggrega- 
tion, so  in  the  process  of  solution,  heat  is  usually  liberated  or  absorbed. 
That  is,  the  solution  is  either  warmer  or  colder  than  the  original 
materials.  This  is  known  as  heat  of  solution.  Thus,  one  gram- 
formula-weight  of  sulphuric  acid  (98  g.),  in  dissolving  in  a  large 
volume  of  water,  liberates  39,170  calories,  and  one  gram-formula- 
weight  of  ammonium  chloride  (53.5  g.),  in  dissolving,  absorbs  3880 
calories. 

The  sources  or  destiny  of  the  heat  given  out  or  absorbed  have  not 
been  studied  in  such  a  way  that  definite  statements  can  be  made  about 
the  theory  of  the  subject.  There  are  many  factors  which  would  have 
to  be  considered.  For  example,  the  body,  if  a  solid,  goes  into  an 
essentially  liquid  condition,  and  its  heat  of  fusion  is  always  negative. 
A  part  of  the  water  always  undergoes  chemical  change  (p.  202).  The 
solute  also  frequently  combines  with  a  part  of  the  water,  or  is  ionized 
(q.v.),  and  the  change  in  volume  of  the  mixture,  as  a  physical 
phenomenon,  would  alone  entail  a  heat-change.  Hence  this  heat 
effect  is  partly  chemical  and  partly  physical  in  origin. 


SOLUTION  203 

The  first  water  used  always  causes  a  greater  heat  change  than  the 
addition  of  succeeding  equal  amounts.  Heats  of  solution  are  meas- 
ured for  the  solution  of  one  formula-weight  of  the  substance  in  un- 
limited water.  The  values  in  calories  for  some  common  substances 
are  as  follows: 

H2SO4,  Aq  =  +  39,170  NaCl,  Aq  =  -1180 

HC1,  Aq  =  +17,400  NH4C1,  Aq  =  -3880 

KOH,  Aq  =  +12,500  KC1,  Aq  =  -4440 

NaOH,  Aq  =  +9780  Na2CO3,10H2O,  Aq  =  -16,160 

Na2CO3,  Aq  =  +5640  Na2SO4,Aq  =  +460 

CaCl2,  Aq  =  +3258  Na2SO4,10H2O,  Aq  =  -18,760 

When  a  substance  comes  out  of  solution,  the  heat  effect  is  equal,  and  of 
opposite  sign  to  that  occurring  when  the  same  substance  goes  into  solu- 
tion. Hence,  since  the  decahydrate  of  sodium  sulphate  absorbs  heat 
in  dissolving,  a  considerable  development  of  heat  is  noticed  when  it 
suddenly  crystallizes  from  a  supersaturated  solution.  Some  ether  in 
a  tube  immersed  in  the  solution  may  be  boiled  by  this  heat  and  its 
vapor  set  on  fire  to  make  the  fact  evident  at  a  distance.  An  impor- 
tant relation  between  heat  of  solution  and  solubility  will  be  discussed 
under  van't  Hoff's  law  of  mobile  equilibrium  (q.v.). 

The  melting  of  ice  by  contact  with  salt  is  sometimes  explained  as 
being  due  to  the  heat  given  out  when  the  salt  dissolves.  The  above 
table  shows,  however,  that  heat  is  absorbed  -M  not  given  out  —  when 
salt  dissolves.  The  fact  is  that  any  substance,  provided  it  is  soluble 
in  water,  when  added  to  ice,  will  give  a  mixture  with  a  temperature 
below  0°.  Whether  the  substance  is  sulphuric  acid,  which  gives  out 
heat,  or  ammonium  chloride,  which  absorbs  it,  makes  no  difference  in 
the  result.  The  solubility,  in  terms  of  the  number  of  molecules  that 
can  be  dissolved  in  a  given  amount  of  water,  alone  determines  the 
extent  of  the  depression  in  temperature  (below  0°)  which  is  produced. 
The  temperature  falls  because  of  the  heat  absorbed  in  melting  the  ice 
(ht.  of  fusion  79  cal.  per  1  g.),  and,  in  spite  of  the  warmth  of  surround- 
ing bodies,  the  temperature  thereafter  remains  at  the  low  value  (so 
long  as  the  ice  holds  out),  because  that  is  the  temperature  at  which 
alone  ice  and  the  solution  can  coexist  (p.  200). 

Definition  of  a  Solution.  —  We  are  now  able  to  make  a  brief 
statement  which  shall  distinguish  solutions  from  mixtures  on  the  one 
hand  and*  from  chemical  compounds  on  the  other.  Solutions  are 
homogeneous  mixtures  of  two  or  more  substances  which  are  separa- 


204  INORGANIC  CHEMISTRY 

ble  into  their  components  by  altering  the  state  of  one  of  the  sub- 
stances (e.g.,  by  freezing  or  boiling  out  one  component),  and  whose 
properties  vary  continuously  with  the  proportions  of  the  components 
between  certain  limits. 

Application  in  Chemical  Work.  —  The  theory  of  this  subject 
has  been  given  on  account  of  its  intensely  practical  interest,  and  it 
should  be  kept  in  mind  in  all  ordinary  chemical  operations.  It  will 
afford  an  explanation  of  many  things  which  might  otherwise  be  attrib- 
uted to  the  wrong  cause,  or  might  remain  entirely  without  explanation. 
For  example,  why  is  the  action  of  a  metal  upon  an  acid  so  slow?  We 
must  remember  that  an  acid  diluted  with  water  is  being  used,  and 
only  one  molecule  out  of  every  dozen  or  hundred  is  a  molecule  of  the 
acid.  So  that  the  access  of  the  latter  to  the  metal  is  restricted  at  first, 
and  becomes  more  and  more  so  as  the  molecules  of  the  acid  in  the 
immediate  vicinity  of  the  metal  undergo  chemical  change.*  On  the 
other  hand,  the  metal,  especially  if  it  be  in  the  form  of  sticks  obtained 
by  casting,  presents  one  of  the  elements  in  the  action  in  a  most  com- 
pact form.  The  only  parts  which  are  accessible  to  the  acid  are  those 
upon  the  surface  and,  the  metal  not  being  appreciably  soluble  in  water, 
the  molecules  can  only  pass  off  and  expose  a  fresh  layer  very  slowly. 
It  is  no  wonder  that  many  chemical  actions  occupy  a  considerable 
time.  The  wonder  is  that  they  should  take  place  as  rapidly  as  they 
do.  Their  speed  would  seem  to  point  to  a  most  intense  chemical 
activity,  even  in  the  seemingly  feebler  instances.  Various  artifices  are 
habitually  employed  for  facilitating  chemical  action.  Thus,  the  metal 
may  be  reduced  to  a  leafy  form  by  pouring  the  molten  substance  into 
cold  water.  Naturally,  with  metals,  the  maximum  surface  and  the 
most  rapid  chemical  action  are  obtained  by  using  a  fine  powder. 

The  most  speedy  interaction  of  all,  other  things  being  equal,  must 
be  attainable  by  dissolving  all  the  interacting  substances  in  water. 
Under  these  circumstances  all  the  molecules  of  each  substance  must 
simultaneously  have  many  molecules  of  the  other  within  easy  reach  of 
them. 

*  In  spite  of  the  continuous  exhaustion  of  the  acid,  there  is  often  a  steady  in- 
crease in  the  rate  at  which  a  dilute  acid  interacts  with  a  metal.  This  is  due,  at 
first,  to  the  dirt  on  the  surface  of  the  metal  which  temporarily  obstructs  the  ac- 
tion and,  later,  to  the  rising  temperature  of  the  interacting  bodies.  In  the  action 
of  metals,  like  that  of  zinc  on  an  acid  (p.  119),  impurities  in  the  metal  remain  on 
the  surface  as  the  metal  is  eaten  away  and  act  as  contact  agents,  stimulating  the 
reaction. 


SOLUTION  205 

Exercises.  —  1.   Give  other  examples  of  limited  solubility  in 
various  solvents  (p.  179). 

2.  If  you  were  not  permitted  to  evaporate  sea-water  to  dryness, 
how  should  you  show  that  it  was  a  solution  and  not  a  pure  substance? 

3.  Reexpress  Henry's  law  (p.  188)  in  terms  of  the  volume  of  gas 
dissolved  at  different  pressures. 

4.  If  hydrogen  sulphide  is  diluted  to  ten  times  its  volume  with 
hydrogen,  what  volume  of  it,  estimated  as  pure  r;as,  will  be  dissolved 
by  20  volumes  of  alcohol  at  0°  and  760  mm.  (p.  188)? 

'5.  If  the  dissolved  air,  after  being  removed  from  water  by  boiling, 
were  to  be  shaken  with  water  once  more,  in  what  proportions  by 
volume  would  the  gases  now  dissolve  (p.  189)? 

6.  Read  from  the  curves  (p.  191)  the  solubilities  of  potassium 
nitrate  at  15°,  of  potassium  chloride  at  30°,  of  potassium  chlorate  at 
45°.     What  are  the  relative  rates  at  which  the  solubilities  of  these 
salts  increase  with  rise  in  temperature? 

7.  Express  the  concentrations  of  solutions  of  ammonium  chloride, 
saturated  at  0°  (sp.  gr.  1.076),  and  of  potassium  sulphate  K2S04,  satu- 
rated at  10°  (sp.  gr.  1.083),  in  terms  of  a  normal  solution  (p.  182). 

8.  Express  the  concentration  of  a  five  per  cent  aqueous  solution 
of  phosphoric  acid  (sp.  gr.  1.027),  in  terms  of  a  normal  and  a  molar 
solution,  respectively. 

9.  Name  the  phases  (p.  192)  in  a  system  consisting  of  oxygen  and 
its  aqueous  solution,  (a)  above  0°,  (6)  below  0°. 

10.  When  a  solution  of  a  very  soluble  substance,  like  zinc  chloride, 
is  evaporated  to  dryness  on  a  water  bath,  why  is  the  escape  of  the 
last  portions  of  the  solvent  so  much  slower  than  is  that  of  the  first? 

11.  At  what  level  in  a  tank  of  water  should  you  introduce  am- 
monia gas,  in  order,  with  the  least  effort,  to  saturate  the  water 
(p.  201)? 

12.  Give  one  reason  why  a  gas,  at  constant  pressure,  should  be- 
come less  soluble  as  the  temperature  rises. 


CHAPTER  XI 


HYDROGEN  CHLORIDE  AND  CHLORINE 

WE  have  had  occasion  several  times  to  mention  common  salt,  or 
sodium  chloride  NaCl.  This  is  one  of  the  most  familiar  chemical 
substances.  Large  quantities  of  it  are  used  in  the  household,  in 
cooking  and  in  making  freezing  mixtures.  Still  larger  amounts  are 
consumed  in  manufacturing  washing  soda,  caustic  soda,  and  soap,  for 
all  of  which  it  furnishes  the  necessary  sodium.  It 
is  used  also  in  preserving  fish  and  other  foods. 
It  supplies  the  chlorine  used  in  bleaching  and  in 
the  sterilization  of  city  waters.  We  shall  consider 
it  first  as  a  means  of  making  other  compounds  of 
chlorine. 


Preparation  of  Hydrogen  Chloride  HCl 
from  Salt.  —  When  some  concentrated  sulphuric 
acid  is  poured  upon  sodium  chloride,  a  vigorous 
effervescence  is  noticed.  This  shows  that  bubbles 
of  a  gas  are  forming  upon  the  salt  crystals  and  are 
rising  through  the  acid  and  bursting.  If  the  salt 
be  placed  in  a  flask  (Fig.  85),  the  acid  can  be 
allowed  to  enter  from  time  to  time  through  the 
funnel.  When  the  air  has  been  displaced  from  the  flask,  the  gas 
issues  from  the  delivery  tube.  If  the  correct  proportion  of  the  acid 
is  used,  and  only  a  gentle  heat  is  applied,  all  that  remains  in  the 
flask  is  a  white  solid,  sodium-hydrogen  sulphate  (or  sodium  bi- 
sulphate)  NaHS04: 


FIG.  85. 


NaCl  +  H2S04  ->  NaHSCVH-  HCl  T-* 


(1) 


The  gas  is  extremely  soluble  in  water  and,  being  heavier  than  air, 
may  be  collected  by  upward  displacement  of  the  air  in  a  jar. 

The  action  described  is  the  one  which  occurs  in  the  laboratory. 

*  An  arrow  pointing  upward  indicates  escape  as  a  gas,  or  solution  of  a  solid; 
one  pointing  downward  indicates  precipitation. 

-    206 


HYDROGEN   CHLORIDE  AND  CHLORINE  207 

When  a  double  proportion  of  salt  and  a  red  heat  are  used,  a  second 
action  occurs: 

NaCl  +  NaHS04  -»  Na2SO4  +  HC1  .  T  (2) 

and  sodium  sulphate  Na^SCX  remains.  In  one  or  two  factories  in 
Europe  this  action  is  still  employed,  with  furnace  heat,  in  manufac- 
turing sodium  sulphate,  from  which  sodium  carbonate  is  afterwards 
prepared.  The  hydrogen  chloride  passes  into  a  tower,  down  which 
water  trickles  over  lumps  of  coke,  and  is  dissolved.  The  aqueous 
solution  is  called  hydrochloric  acid  or,  in  commerce,  muriatic  acid 
(Lat.  muria,  brine). 

Hydrogen  Chloride  from  Other  Chlorides  and  Other  Acids. 

—  The  chlorides  of  other  metals  can  be  substituted  for  sodium 
chloride  in  this  action,  and  all  the  more  soluble  ones  give  hydrogen 
chloride  freely.  Other  chlorides  are  all  more  expensive,  however, 
than  common  salt. 

All  acids  contain  the  necessary  hydrogen  radical,  and  might  offer 
it  in  exchange  for  the  sodium  in  the  salt,  yet  in  practice  no  other  acid 
works  so  well  as  does  sulphuric  acid.  Most  other  acids  contain  much 
water,  which  dissolves  the  hydrogen  chloride  (see  below).  Concen- 
trated phosphoric  acid  H3PO4,  Aq  acts  slowly,  giving  primary  sodium 
phosphate : 

NaCl  +  H3P04  ->  NaH2P04  +  HC1  T  - 

If  various  chlorides  are  used  with  the  same  acid,  it  will  be  found 
that  the  vigor  of  the  actions  is  very  different.  In  some,  hydrogen 
chloride  will  be  produced  copiously  without  the  assistance  of  heat. 
In  others,  there  will  be  difficulty  in  showing  that  hydrogen  chloride 
gas  is  produced  at  all.  We  must  not  hastily  assume  that  this  is  owing 
to  any  greater  chemical  affinity  in  one  case  than  another.  More  ex- 
tensive experimentation  will  show  that  the  more  soluble  chlorides  as  a 
rule  give  more  vigorous  effects  than  those  which  are  less  so  (pp.  204. 
220).  Ammonium  chloride  with  sulphuric  acid  would  represent  the 
former  variety,  while  mercuric  chloride  with  the  same  acid  would 
represent  the  latter. 

The  Molecular  View  of  the  Interaction  of  Sulphuric  Acid 
and  Salt.  —  One  who  has  used  the  above  method  for  making  hy- 
drogen chloride  without  reflection  would  not  realize  the  complexity 
of  the  machinery  by  which  the  result  is  achieved.  The  means  are 
apparently  very  simple.  Yet  the  mechanical  features  of  this  expert- 


208  INORGANIC  CHEMISTRY 

ment,  when  laid  bare,  are  extremely  curious  and  interesting.  A  single 
fact  will  show  the  possibilities  which  are  concealed  in  it. 

If  we  take  a  saturated  solution  of  sodium-hydrogen  sulphate  (that 
is,  one  containing  the  minimum  amount  of  water),  and  add  to  it  a 
concentrated  solution  of  hydrogen  chloride  in  water  (concentrated 
hydrochloric  acid),  we  shall  perceive  at  once  the  formation  of  a 
copious  precipitate.  This  is  composed  entirely  of  minute  cubes  of 
sodium  chloride: 

NaHS04  +  HC1  -»  H2S04  +  NaCl  |  .*  (3) 

Now  this  action  is  nothing  less  than  the  precise  reverse  of  (1),  yet  it 
proceeds  with  equal  success.  In  fact,  this  chemical  interaction  is  not 
only  reversible  (p.  153),  but  can  be  carried  to  completion  in  either  direc- 
tion. It  is  only  in  presence  of  a  large  amount  of  water,  sufficient  to 
keep  both  the  hydrogen  chloride  and  the  salt  in  solution,  that  it  stops 
midway  in  its  career  and  is  valueless  for  securing  a  complete  trans- 
formation in  either  direction : 

NaHSO4  +  HC1  <±  H2S04  +  NaCl.  (4) 

In  an  action  which  is  reversible,  if  the  products  remain  as  perfectly 
mixed  and  accessible  to  each  other  as  were  the  initial  substances, 
because  all  are  in  solution  (4),  their  interaction  will  continually  undo 
a  part  of  the  work  of  the  forward  direction  of  the  change.  Hence,  in 
such  a  case  the  reaction  must,  and  does,  come  to  a  standstill  while  as 
yet  only  partly  accomplished.  But  this  was  not  the  case  with  actions 
(1)  and  (3).  Let  us  examine  the  means  by  which  the  premature 
cessation  of  each  was  avoided. 

In  (1)  the  salt  dissolved  to  some  extent  in  the  sulphuric  acid, 
NaCl  (solid)  <=±  NaCl  (dslvd),  and  so,  by  contact  of  the  two  kinds  of 
molecules,  the  products  were  formed.  On  the  other  hand,  the  hydro- 
gen chloride,  being  almost  insoluble  in  sulphuric  acid,  escaped  as  fast 
as  it  was  formed:  HC1  (dslvd)  <=^  HC1  (gas).  Hence,  in  that  case, 
almost  no  reverse  action  was  possible,  and  the  double  decomposition 
went  on  to  completion.  With  all  the  sodium-hydrogen  sulphate  in 
the  bottom  of  the  flask,  and  most  of  the  hydrogen  chloride  in  the  space 
above,  the  two  products  might  as  well  have  been  in  separate  vessels 
so  far  as  any  efficient  re-interaction  was  concerned.  This  plan,  in 
which  water  is  purposely  excluded,  forms  therefore  the  method  of 
making  hydrogen  chloride. 

NaCl  (solid)  *±  NaCl  (dslvd)  +  H2SO4  *±  NaHSO4  +  HC1  (dslvd)  «*  HC1  (gas), 
*  See  footnote  on  p.  200. 


HYDROGEN   CHLORIDE  AND  CHLORINE  209 

In  (3),  on  the  other  hand,  the  hydrogen  chloride  was  taken  in 
aqueous  solution,  and  was  mixed  with  a  strong  solution  of  sodium  bi- 
sulphate.  The  acid  was,  therefore,  kept  permanently  in  full  contact 
with  the  sodium  bisulphate.  It  had,  in  this  case,  every  opportunity 
to  interact  with  the  latter  and  no  chance  of  escape.  Every  molecule 
of  each  ingredient  could  reach  every  molecule  of  the  other  with  equal 
ease.  Furthermore,  the  sodium  chloride  produced  as  a  result  of  their 
activity  is  not  very  soluble  in  concentrated  hydrochloric  acid  (far 
less  so  than  in  water),  and  so  it  came  out  as  a  precipitate:  NaCl 
(dslvd)  <=±  NaCl  (solid).  But  this  was  almost  the  same  as  if  it  had 
gone  off  as  a  gas.  It  meant  that  the  greater  part  of  the  salt  was  in  the 
solid  form.  It  was  in  a  state  of  fine  powder,  it  is  true.  But,  in  the 
molecular  point  of  view,  the  smallest  particle  of  a  powder  contains 
millions  of  molecules,  and  most  of  these  are  necessarily  buried  in  the 
interior  of  a  particle.  Thus,  the  sodium  chloride  was  no  longer  able 
to  interact  effectively  molecule  to  molecule  with  the  other  product, 
the  sulphuric  acid.  Hence,  there  was  little  reverse  action  to  impede 
the  progress  of  the  primary  one.  Thus  (3)  is  nearly  as  perfect  a  way 
of  liberating  sulphuric  acid  as  (1)  is  of  liberating  hydrogen  chloride. 

This  discussion  is  given  to  illustrate  the  displacement  of  a  chemi- 
cal equilibrium,  and  to  explain  the  method  of  preparing  hydrogen 
chloride.  It  also  throws  an  interesting  light  on  chemical  affinity, 
however.  Considering  action  (1),  by  itself,  we  might  reason  that  the 
hydrogen  chloride  was  formed  because  the  affinity  of  the  hydro- 
gen (H)  for  chlorine  (Cl)  was  greater  than  for  the  sulphate  radical 
(SO4).  But,  if  we  did  so,  then  in  action  (3)  we  should  be  compelled 
to  reason  similarly  that  the  preponderance  of  affinity  was  just  the 
opposite.  In  point  of  fact,  no  conclusion  about  relative  affinity 
can  be  drawn  from  these  actions.  The  effects  of  affinity  are  here 
entirely  subordinated  by  the  effects  of  a  purely  mechanical  arrange- 
ment (pp.  37,  127,  204),  depending  on  solubility. 

The  egregious  misconception  that  sulphuric  acid  is  shown  by  this  action  to  be 
"stronger"  than  hydrochloric  acid  was  disposed  of,  so  far  as  the  science  was  con- 
cerned, half  a  century  ago.  But  it  survives  in  suburban  chemical  circles  with 
remarkable  tenacity.  The  fact,  quaintly  enough,  is  that  real  relation  in  respect 
to  activity  is  just  the  reverse. 

Other  Ways  of  Obtaining  Hydrogen  Chloride.  —  Although 
never  used  for  generating  hydrogen  chloride  on  a  large  scale,  there  is 
another  important  kind  of  action  in  which  this  substance  is  a  product. 
When  water  acts  upon  the  chlorides  of  non-metallic  substances,  like 


210  INORGANIC  CHEMISTRY 

sulphur,  phosphorus,  and  iodine,  a  double  decomposition  (p.  20) 
occurs.  Since  water  is  always  one  of  the  interacting  substances, 
this  kind  of  change,  —  a  double  decomposition  involving  water,  —  is 
called  hydrolysis  (Gk.  v8<op,  water,  and  XVO-LS,  the  act  of  loosing). 
Thus,  when  a  little  water  is  added  to  one  of  the  chlorides  of  phos- 
phorus, hydrogen  chloride  is  formed.  Besides  this  substance,  the 
trichloride  gives  phosphorous  acid,  and  the  pentachloride,  phosphoric 
acid: 

PC13  +  3HOH  -» 3HC1  +  P(OH)3. 
PC15  +  4H20  -»  5HC1  +  H3P04. 

The  water  divides  into  the  radicals  H  and  OH,  and  the  former 
unites  with  the  more  active  non-metallic  element  in  the  substance 
(the  Cl,  in  PCls)  and  the  hydroxyl  with  the  other  element. 

A  dissociation  is  a  reversible  decomposition  of  one  substance  into  two  or  more 
(p.  148).  Hydrolysis  is  an  ordinary  double  decomposition  or  metathesis  where 
water  is  one  of  the  reagents.  Yet  it  has  been  perversely  named  hydrolytic  dis- 
sociation by  many  writers.  A  whole  chapter  might  be  devoted  to  the  ingenuity 
with  which  chemists  have  misnamed  many  of  the  things  with  which  they  deal. 
Perhaps  this  tendency  is  a  survival  of  the  habit  the  alchemists  had  of  using 
obscure  and  symbolical  names  for  their  materials  to  prevent  the  penetration  of 
their  secrets  by  uninitiated  seekers  after  knowledge.  Important  facts  and  prin- 
ciples have  been  sedulously  labeled  with  misleading  titles,  like:  Water  of  crystalli- 
zation, which  has  no  more  to  do  with  crystallization  than  with  color,  density,  or 
any  other  physical  property;  supersaturated  solution,  which,  as  a  solution,  is  the 
same  as  any  other;  mass  action,  which  has  nothing  to  do  with  mass,  but  is  con- 
cerned wholly  with  concentration;  strong  acid,  which  refers  to  activity  and  not 
power  of  resistance;  reciprocal  proportions,  a  law  in  which  reciprocals  of  numbers 
play  no  part;  downward  displacement  of  air,  when  the  air  is  displaced  upwards, 
and  so  forth.  Here  there  is  an  opportunity  to  confuse  hydrolysis  with  electrolytic 
dissociation,  and  the  beginner  never  fails  to  embrace  it.  Hydrolytic  double  de- 
composition would  have  been  a  correct,  if  somewhat  clumsy  term. 

Often,  when  a  steady  stream  of  hydrogen  chloride  is  required,  con- 
centrated hydrochloric  acid  is  placed  in  a  generating  flask,  and  concen- 
trated sulphuric  acid  is  allowed  to  trickle  into  it  from  a  dropping 
funnel.  The  hydrogen  chloride  is  less  soluble  in  diluted  sulphuric 
acid  than  in  water  (see  Solubility  product)  and  escapes. 

Physical  Properties.  —  Hydrogen  chloride  is  a  colorless  gas, 
which  produces  a  suffocating  effect  when  breathed. 

Density  (H  =  1),  18.23  Grit,  temp.,  +52° 

Weight  of  22.4  1.,  36.73  g.  Boiling-point  (liq.),  -83.7° 

Solubility  in  Aq.  (0°),  50,300  vols.  in  100    Melting-point  (solid),  -110° 


HYDROGEN  CHLORIDE  AND  CHLORINE 


211 


FIG.  86. 


The  gas  is  one-fourth  heavier  than  air.  On  account  of  its  great 
solubility  and  the  small  vapor  tension  of  its  solution,  it  condenses  at- 
mospheric moisture  into  a  fog  of  drops  of  hydrochloric  acid.  Its 
extreme  solubility  may  be  shown  by  filling  a  dry  flask  (Fig.  86)  with 
the  gas.  The  "  dropper  "  contains  water,  and  is  closed 
at  the  tip  with  soft  wax.  A  drop  of  water,  expelled  by 
pinching  the  "  dropper,"  dissolves  so  much  of  the  gas 
that  the  water  is  forced  in  by  atmospheric  pressure, 
like  a  fountain,  through  the  longer  tube.  On  account 
of  its  high  critical  point,  it  may  be  liquefied  by  pressure 
alone.  Both  in  the  gaseous  and  liquefied  states  it  is  a 
nonconductor  of  electricity.  .Its  heat  of  solution  (p. 
203)  is  17,400  calories. 

On  account  of  its  high  concentration,  the  solution 
may  be  looked  upon  as  a  mixture  of  liquefied  hydrogen 
chloride  and  water.  At  15°  and  760  mm.  454.6  volumes 
of  the  gas  dissolve  in  1  volume  of  water,  or  746  g.  in  1  1. 
The  mixture  weighs  therefore  1746  g.  (42.7  per  cent  of 
HC1).  Its  sp.  gr.  is  1.215.  The  volume  of  the  solution 
is  given  by  the  proportion  1215  : 1  :  :  1746  :  x,  in  which  x  =  1.437  1. 
Hence  the  addition  of  454.6  liters  of  the  gas  has  increased  the  volume 
by  only  437  c.c.  Now  at  15°  the  sp.  gr.  of  liquefied  hydrogen  chlo- 
ride is  0.8320,  and  the  volume  of  746  g.  is  therefore  746  -r-  0.832  = 
896  c.c.  So  that,  even  if  the  substances  had  been  mixed  in  liquid 
form,  a  considerable  shrinkage  would  still  have  occurred  (1000  c.c. 
Aq  +  896  c.c.  liq.  HC1  ->  1437  c.c.  HC1,  Aq). 

When  the  concentrated  aqueous  solution  is  heated,  it  is  the  gas  and 
not  the  water  which  is  driven  out,  for  the  most  part.  When  the  con- 
centration has  been  reduced  to  20.2  per  cent  the  rest  of  the  mixture 
distils  unchanged  at  110°.  This  occurs  because,  at  this  concentra- 
tion, the  gas  enters  into,  and  is  carried  off  in,  the  bubbles  of  steam  in 
the  same  proportion  in  which  it  is  present  in  the  liquid.  With  a 
concentration  above  20.2  per  cent,  more  hydrogen  chloride  enters  the 
bubbles,  with  one  below  20.2  per  cent,  more  water.  Hence,  if  a 
dilute  solution  is  used,  water  is  the  chief  product  of  distillation  (about 
100°),  but  gradually  the  boiling-point  rises  and,  when  the  concentra- 
tion has  reached  20.2  per  cent  once  more,  the  same  hydrochloric  acid 
of  constant  boiling-point  (110°  at  760  mm.),  as  it  is  called,  forms  the 
residue.  It  is  thus  impossible  to  separate  by  distillation  the  com- 
ponents of  mixtures  which  behave  in  this  way.  This  must  necessarily 
be  the  case  whenever,  as  here,  the  vapor  tensions  of  the  components 


212  INORGANIC   CHEMISTRY 

separately  and  those  of  all  other  mixtures  are  higher  (and  the  b.-ps. 
lower)  than  that  of  one  particular  mixture.  When,  as  is  more  often 
the  case,  one  of  the  components  has  a  vapor  tension  which  is  lower 
than  that  of  the  other  and  lower  than  that  of  any  mixture  of  the  two, 
this  component  will  tend  to  remain  behind,  and  separation  can  be 
effected.  The  separation  of  petroleum  products  (q.v.)  from  one  an- 
other illustrates  the  common  case  (see  under  Alcohol  for  the  third 
possibility). 

The  composition  of  the  mixture  having  the  minimum  vapor  ten- 
sion varies  with  the  external  pressure,  and  so  does  the  boiling-point. 
At  300  mm.  the  constant  boiling  liquid  contains  21.8  per  cent  of 
hydrogen  chloride  and  boils  at  84°;  at  1520  mm.  it  contains  19.1  per 
cent  of  the  gas. 

The  common  belief  that  hydrochloric  acid  of  constant  boiling-point  is  a  defi- 
nite compound  is  without  foundation.  Compounds  do  not  vary  in  composition 
with  changes  in  pressure  in  this  manner.  Aqueous  solutions  of  hydrogen  iodide, 
hydrogen  bromide,  and  nitric  acid  behave. in  the  same  way.  But  solutions  of 
oxygen,  of  ammonia,  and  of  many  liquids  (e.g.,  methyl  alcohol)  in  water  belong 
to  the  second  of  the  two  classes  mentioned  above,  and  the  more  volatile  com- 
ponent often  leaves  the  water  entirely  before  much  of  the  latter  has  evaporated. 

Chemical  Properties.  —  Hydrogen  chloride  is  extremely  stable, 
as  we  might  expect  from  the  vigor  with  which  the  elements  of  which  it 
is  composed  combine  (see  p.  222).  On  being  heated  to  a  temperature 
of  1800°  it  dissociates  into  its  constituents  to  a  slight  extent. 

In  the  chemical  point  of  view,  it  is  on  the  whole  rather  an  indif- 
ferent substance.  When  water  is  saturated  with  the  gas  at  —  22°  a 
hydrate  HC1,2H2O  crystallizes  out.  This  decomposes  into  the  same 
constituents  when  allowed  to  warm  up  again  to  —18°.  Hydrogen 
chloride  (the  gas)  has  no  action  upon  any  of  the  non-metals,  such  as 
phosphorus,  carbon,  sulphur,  etc.  Many  of  the  metals,  however,  par- 
ticularly the  more  active  ones,  such  as  potassium,  sodium,  and  mag- 
nesium, decompose  it.  Hydrogen  is  set  free,  and  the  chloride  of 
the  metal  is  formed  (K  +  HC1 ->  KC1  +  H).  Hydrogen  chloride 
unites  directly  with  ammonia  gas  to  form  a  cloud  of  solid  particles  of 
ammonium  chloride  (HC1  +  NH3  — »  NH4C1).  The  liquefied  gas  has 
the  same  properties. 

Chemical  Properties  of  Hydrochloric  Acid.  —  The  solution 
of  hydrogen  chloride  in  water  is  an  entirely  different  substance  in  its 
chemical  behavior  from  hydrogen  chloride,  It  is  strongly  acid,  turn- 


HYDROGEN   CHLORIDE  AND  CHLORINE  213 

ing  blue  litmus  red.  The  gas  and  liquefied  gas  have  no  such  property. 
The  solution  conducts  electricity,  as  we  have  seen  (p.  120),  very  well, 
and  is  decomposed  in  the  process,  giving  hydrogen  at  the  negative 
wire  and  chlorine  at  the  positive  wire: 

HC1  -»  H  (neg.  wire)  +  Cl  (pos.  wire). 

The  gas  and  the  liquefied  gas  are  practically  nonconductors. 

The  metals  preceding  hydrogen  in  the  order  of  activity  (p.  129), 
when  introduced  into  hydrochloric  acid,  displace  the  hydrogen  (p. 
118),  and  form  the  chloride  of  the  metal.  In  the  case  of  zinc  the 
action  was  represented  by  the  equation: 


The  liquefied  gas  has  no  action  upon  zinc,  and  even  its  solutions  in 
many  other  solvents  show  little  activity.  The  solution  in  alcohol 
behaves  like  that  in  water.  But  the  solutions  in  toluene,  benzene,  and 
other  compounds  of  carbon  and  hydrogen,  in  many  of  which  the  gas  is 
freely  soluble,  are  hardly  affected  by  the  presence  of  zinc  and  other 
metals.  These,  and  many  other  facts  which  we  shall  notice  later 
(see  Dissociation  in  solution),  show  that  the  condition  of  this  sub- 
stance in  aqueous  solution  is  peculiar. 

The  aqueous  solution  of  hydrogen  chloride  interacts  rapidly  with 
most  oxides  and  hydroxides  of  metals,  as,  for  example,  those  of  zinc: 

ZnO  +  2HC1  -»  ZnCl2  +  H2O. 
Zn(OH)2  +  2HC1  ->  ZnCl2  +  2H2O. 

Here  no  hydrogen  is  obtained,  since  the  oxygen  in  the  oxide,  and 
the  hydroxyl  in  the  hydroxide,  unite  with  it  to  form  water.  In  each 
case,  however,  the  chloride  of  the  metal  is  obtained.  It  may  be  noted 
in  passing  that  all  acids  behave  in  a  similar  manner  towards  oxides 
and  hydroxides,  giving  water  and  a  compound  corresponding  to  the 
chloride  (see  p.  218).  Dilute  sulphuric  acid,  for  example,  gives 
sulphates. 

Modes  of  Preparing  Chlorides.  —  In  the  preceding  section 
three  kinds  of  actions,  each  constituting  a  different  way  of  obtaining 
chlorides,  have  been  mentioned  incidentally.  There  are  two  others. 
The  simplest  is  the  direct  union  of  the  element  with  chlorine  (Zn  -f 
2C1—  >ZnCl2).  The  other  method  is  illustrated  in  the  case  of  the 
precipitation  of  silver  chloride  (p.  20),  by  adding  a  solution  of  a 


214  INORGANIC   CHEMISTRY 

chloride  to  a  solution  of  silver  nitrate.  Here  the  formation  of  the 
chloride  occurred  by  exchange  of  another  radical  for  chlorine: 

AgN03  +  NaCl  ->  AgCl  |  +  NaNO3. 

The  insoluble  chlorides  (see  p.  226)  can  be  made  conveniently  by  this 
plan.  The  formation  of  the  precipitates,  for  example  that  of  silver 
chloride,  is  used  as  a  test  for  the  presence  of  a  soluble  chloride  in  the 
solution. 

Double  decompositions,  like  the  action  just  mentioned,  involving 
acids,  bases,  and  salts  (see  below),  are  all  reversible  reactions.  The 
fact  that  many  of  them  proceed,  nevertheless,  to  practical  completion 
has  already  been  explained  at  length  (p.  208). 

Uses  of  Hydrochloric  Acid.  — •  This  substance  is  employed  in 
cleaning  metals,  and  in  the  manufacture  of  chlorides  of  metals.  It 
is  an  important  component  of  the  gastric  juice  of  the  stomach, 
although  the  proportion  is  only  about  1  part  in  500. 

Precipitation.  —  When  two  soluble  substances  are  dissolved, 
separately,  and  the  solutions  are  mixed,  chemical  interaction  fre- 
quently occurs,  as  in  the  case  of  salt  and  silver  nitrate  (see  also  p. 
209).  If  one  of  the  products  is  insoluble,  then  a  supersaturated 
solution  of  this  product  is  at  once  produced.  As  a  rule,  this  sub- 
stance almost  immediately  becomes  visible  as  a  fine  powder,  called  a 
precipitate,  suspended  in  the  liquid. 

The  insoluble  product  can  often  be  recognized  by  its  physical 
appearance,  and  so  this  sort  of  action  is  frequently  used  as  a  test  for 
one  of  the  original  substances.  Thus  precipitates  have  often  a 
distinctive  color.  Again,  precipitates  which  are  colorless,  or  have  the 
same  color,  differ  in  appearance,  and  are  described  as  gelatinous, 
curdy,  pulverulent,  or  crystalline.  In  the  first  two  cases,  the  precipi- 
tation is  so  sudden  that  there  is  not  time  for  crystals  to  be  formed, 
and  the  product  is  amorphous  (Gk.  a,  priv.;  ^op<j>^  form).  Thus 
silver  chloride  is  curdy,  and  precipitated  sodium  chloride  (p.  208)  is 
crystalline. 

Salts.  —  We  have  seen  that  an  acid  contains  hydrogen  as  a 
radical  (p.  117),  and  a  base  contains  the  radical  hydroxyl  OH  (p.  149). 
The  name  salts  is  given  to  the  class  of  substances,  which  contain  a 
positive  and  a  negative  radical,  neither  of  which  is  hydrogen  or 
hydroxyl.  For  example,  NaCl,  Na2S04,  AgNOa  are  the  formulae  of 


HYDROGEN   CHLORIDE  AND  CHLORINE  215 

salts.  They  are  so  named  because  they  resemble  common  salt,  in 
having  two  radicals,  and  entering  readily  into  double  decompo- 
sition. 

Sodium-hydrogen  sulphate  NaHSC>4  is  classed  as  an  acid  salt, 
because  it  has  a  positive  and  a  negative  radical,  and  a  hydrogen 
radical  in  addition. 

CHLORINE 

Chlorine  was  first  recognized  as  a  distinct  substance  by  Scheele 
(1774).  He  obtained  it  from  salt  by  means  of  manganese  dioxide, 
using  the  common  method  described  below.  It  was  for  years  sup- 
posed to  be  a  compound  containing  oxygen,  until  Davy  (1809-1818) 
demonstrated  that  it  was  an  element. 

Occurrence.  —  Chlorine  does  not  occur  free  in  nature.  There 
are,  however,  many  compounds  of  it  to  be  found  in  the  mineral 
kingdom.  Sea- water  contains  a  number  of  chlorides  in  solution. 
Nearly  2.8  of  the  3.6  per  cent  of  solid  matter  in  sea-water  is  sodium 
chloride  NaCl.  During  past  geological  ages  the  evaporation  of  sea- 
water  has  led  to  the  formation  of  immense  deposits  of  the  compounds 
usually  found  in  such  water.  Thus,  at  Stassfurt,  such  strata  attain  a 
thickness  of  over  a  thousand  feet.  Certain  layers  of  these  strata  are 
composed  mainly  of  sodium  chloride,  called  by  the  mineralogist  halite 
(rock  salt).  In  other  layers  potassium  chloride  (sylvite),  an  indis- 
pensable fertilizer,  and  hydrated  magnesium  chloride  (bischofite), 
and  other  compounds  of  chlorine,  occur. 

Preparation.  —  Chlorine  cannot  be  obtained  with  the  same 
ease  as  oxygen.  There  are  only  a  few  chlorides,  such  as  those  of  gold 
and  platinum,  which  lose  chlorine  when  heated,  and  they  are  too 
expensive  or  difficult  to  make  for  laboratory  use.  We  employ  there- 
fore methods  like  those  used  for  the  preparation  of  hydrogen  (cf. 
p.  122).  We  may  (1)  decompose  any  chloride  by  means  of  electricity, 
just  as,  to  get  hydrogen,  we  electrolyzed  a  dilute  acid  (p.  120).  Or  (2) 
we  may  take  some  inexpensive  compound  of  chlorine,  such  as  hydro- 
gen chloride  HC1,  and  by  means  of  some  simple  substance  which  is 
capable  of  uniting  with  the  hydrogen,  —  here  oxygen  serves  the 
purpose,  —  secure  the  liberation  of  the  chlorine.  Or  (3)  —  and  this 
turns  out  to  be  the  most  convenient  laboratory  method  —  we  may 
use  a  more  complex  action. 


216 


INORGANIC   CHEMISTRY 


FIG.  87. 


Electrolysis  of  Chlorides.  —  Hydrogen  chloride  and  those  chlo- 
rides of  metals  which  are  soluble  in  water  are  all  decomposed  when  a 
current  of  electricity  is  passed  through  the  aqueous  solution.  They 
yield  chlorine  at  the  positive  electrode.  The  other  constituent,  the  hy- 
drogen (Fig.  87),  manganese,  or  whatever  it  may  be,  migrates  towards 

the  negative  wire.  To  decompose 
hydrochloric  acid  an  electromotive 
force  of  at  least  1.31  volts  is  re- 
quired. Since  the  chlorine  is  solu- 
ble in  water,  the  effervescence  due 
to  its  release  is  not  noticeable  un- 
til the  liquid  round  the  electrode 
has  become  saturated  with  the  gas : 
Cl  (dslvd)  <=»  Cl  (gas).  The  shape 
of  the  apparatus  keeps  the  two 
products  from  mingling.  The 
presence  of  the  chlorine  in  the 
liquid  at  the  positive  end  may  be 
shown  by  a  suitable  test  (pp.  89 
and  223). 

In  commerce  chlorine  is  now 
obtained  chiefly  by  this  method,  sodium  chloride  or  potassium 
chloride  being  the  source  of  the  element.  Electrodes  of  artificial 
graphite  are  used,  as  most  other  conductors  unite  with  the  chlorine. 
The  potassium  or  sodium,  as  the  case  may  be,  travels  towards  the 
negative  electrode,  but  is  not  liberated.  Instead,  potassium  or  so- 
dium hydroxide  (q.v.)  accumulates  in  the  solution  round  the  plate  and 
hydrogen  escapes.  The  chlorine  is  released  at  the  positive  electrode,  as 
usual.  The  hydrogen,  the  hydroxide  and  the  chlorine  all  find  commer- 
cial applications.  The  chlorine  is  either  liquefied  by  compression  in 
iron  cylinders  or  employed  at  once  for  making  bleaching  powder  (q.v.). 

Action  of  Free  Oxygen  on  Chlorides.  —  Sodium  chloride  is 
the  cheapest  source  of  chlorine,  but  oxygen  does  not  interact  with  it 
even  at  a  high  temperature.  Hence  the  chlorine  must  first  be 
transferred  to  some  other  form  of  combination.  By  treating  the 
sodium  chloride  with  sulphuric  acid,  therefore,  the  chlorine  is  first 
transferred  into  combination  with  the  hydrogen  of  the  acid,  giving 
hydrogen  chloride  (p.  206).  In  order  to  liberate  chlorine  from  this 
compound,  we  may  combine  the  hydrogen  with  oxygen  obtained  from 
the  air: 

2HCH-O«=±H*0 


HYDROGEN   CHLORIDE  AND  CHLORINE  217 

The  two  gases  interact  so  slowly,  however,  that  a  contact  agent  must 
be  employed.  The  mixture  of  air  and  hydrogen  chloride  is  passed 
over  pieces  of  heated  pumice-stone  or  broken  brick,  previously 
saturated  with  cupric  chloride  solution  (Fig.  88) .  A  temperature  of 
about  370°  is  used.  Furthermore,  the  action  is  reversible  (read  the 
equation  backwards)  and  equilib- 
rium is  reached  when  80  per  cent 
of  the  hydrogen  chloride  has  been  FlQ  ^ 

decomposed.     Hence  20  per  cent 

of  this  gas  passes  on  unchanged.  Only  80  per  cent  of  the  hydrogen 
chloride  and  oxygen  are  changed  into  steam  and  chlorine,  because  the 
latter  substances  are  continuously  interacting  to  reproduce  hydrogen 
chloride  and  oxygen.  If  one  substance  (the  steam  or  the  chlorine) 
could  be  separated  from  the  other  (p.  208),  to  prevent  the  backward 
action,  the  yield  would  be  raised  to  100  per  cent.  But  even  a  very 
partial  separation  of  two  gases  requires  elaborate  apparatus  and 
complete  separation  is  practically  impossible.  In  the  product,  there- 
fore, the  chlorine  is  mixed  with  steam,  as  well  as  with  a  very  large 
volume  of  nitrogen  which  entered  with  the  oxygen,  so  that  for  making 
the  pure  substance  this  method  (Deacon's  process)  is  quite  unsuitable. 
Bleaching  powder  (q.v.),  however,  can  be  made  by  this  process. 

Using  the  same  principle,  magnesium  chloride  may  be  heated  in  a  stream  of  air, 
when  the  oxide  of  magnesium  is  formed  and  chlorine  is  given  off :  MgCl2  +  O  — > 
MgO  +  2C1.  The  oxide  of  magnesium  can  then  be  treated  with  hydrochloric 
acid  to  regenerate  the  chloride,  which  in  turn  may  be  subjected  once  more  to  the 
action  of  oxygen.  The  process  is  thus  a  continuous  one. 

The  above  action  is  spoken  of  as  an  oxidation.  It  is  true  that  no 
oxygen  is  actually  introduced  into  the  hydrogen  chloride  as  a  whole. 
The  removal  of  hydrogen  from  combination  with  the  chlorine  is,  how- 
ever, the  first  step  towards  the  introduction  of  oxygen  into  combina- 
tion with  the  latter,  which  is  essentially  an  oxidation. 

Action  of  Combined  Oxygen  upon  Chlorides.  —  The  best 
laboratory  method  for  making  chlorine  is  to  place  some  solid  potas- 
sium permanganate  in  a  flask,  arranged  like  that  in  Fig.  89.  Con- 
centrated hydrochloric  acid  (an  aqueous  solution  of  hydrogen  chloride) , 
diluted  with  one-third  of  its  volume  of  water,  is  allowed  to  fall  upon 
the  compound  drop  by  drop  from  the  dropping  funnel.  The  action 
is  very  rapid,  the  acid  is  exhausted  almost  as  fast  as  it  falls,  and  so  the 
stream  of  gas  can  be  stopped  by  simply  closing  the  stopcock.  The 


218 


INORGANIC  CHEMISTRY 


FIG.  89. 


gas  is  passed  through  a  washing  bottle  containing  water,  in  order  to 
remove  any  hydrogen  chloride  which  may  be  carried  over.  It  may 
be  clried,  if  necessary,  in  a  second  washing  bottle  containing  concen- 
trated sulphuric  acid.  It  cannot 
be  collected  over  water  on  ac- 
count of  its  solubility,  so  that 
jars  are  usually  filled  with  it  by 
upward  displacement  of  air. 

Skeleton:   KMn04-  +  HC1  ->  H20 
+  KC1  +  MnCl2  +  Cl. 

The  O4,  being  all  converted  into 
water,  requires  8H,  and  therefore 
8HC1,  for  the  action.  The  two 
metals,  potassium  and  manga- 
nese, give  their  respective  chlo- 
rides, KC1  and  MnCl2.  This  uses 
3C1,  and  hence  5C1  remains  over 
to  be  liberated: 

Balanced:    KMnO4  +  8HC1  -*  4H20  +  KC1  +  MnCl2  +  5C1. 

The  combined  oxygen  of  the  permanganate  has  oxidized  the  hydro- 
gen chloride,  just  as  did  the  free  oxygen  in  Deacon's  process. 

Other  Means  of  Oxidizing  Hydrogen  Chloride.  —  Many 
other  compounds  of  oxygen  with  metals  interact  with  hydrochloric 
acid  to  give  free  chlorine.  Lead  dioxide  Pb02,  potassium  chlorate 
KClOs,  potassium  dichromate  K2Cr2C>7,  and  manganese  dioxide 
MnO2,  are  of  this  nature.  The  last,  being  inexpensive,  is  commonly 
used  in  making  chlorine.  Being  an  insoluble  substance,  however, 
the  manganese  dioxide  acts  much  more  slowly  than  does  the  potassium 
permanganate,  which  is  soluble.  A  large  amount  of  the  materials, 
and  the  aid  of  heat,  are  required  to  secure  a  rapid  stream  of  chlorine. 

Manganese  Dioxide  and  Hydrogen  Chloride.  —  The  action 
of  manganese  dioxide  upon  hydrochloric  acid  is  an  instructive  one. 
It  is  a  general  rule,  of  which  we  shall  meet  many  applications,  that 
when  an  acid  interacts  with  an  oxide  of  a  metal,  there  are  two  con- 
stant features  in  the  result,  namely:  (1)  The  oxygen  of  the  oxide 
combines  with  the  hydrogen  of  the  acid  to  form  water,  and  (2) 
the  metal  of  the  oxide  combines  with  the  acid  radical  of  the  acid 
according  to  the  valences  of  each.  Here  the  skeleton  equation  should 


HYDROGEN   CHLORIDE  AND  CHLORINE  219 

be  Mn02  +  HC1  m  H20  +  MnCl4.  With  02,  to  form  water,  4HC1  is 
required,  and  the  product  is  2H2O.  Hence  the  equation  is 

Balanced:  Mn02  +  4HC1  ->  2H20  +  MnCU. 

This  is  what  happens  in  the  first  place.  The  products  actually 
obtained,  however,  are  water,  manganous  chloride  MnCl2,  and 
chlorine.  The  manganese  tetrachloride  is  decomposed  by  the 
heating,  the  chlorine  escapes,  and  the  other  two  products  remain  in 
the  vessel. 

MnO2  +  4HC1  ->  2H2O  +  MnCl2  +  2C1.  (1) 

We  owe  the  chlorine  to  the  fact  that  the  tetrachloride  is  unstable. 

When  the  mixture  is  surrounded  by  ice  and  saturated  with  chlorine,  it  can  be 
shown  that  it  contains  the  tetrachloride.  If  it  is  quickly  poured  into  water, 
hydrated  manganese  dioxide  is  precipitated  (Wacker).  The  decomposition  of 
the  tetrachloride  is  reversible: 

MnCU  ±H»  MnCl2  +  2C1, 

and  is  driven  back  by  the  excess  of  chlorine.  The  tetrachloride  is  hydrolyzed  by 
water: 

MnCU  +  2H2O  +  zH2O  —  »  MnO2,zH2O  +  4HC1. 


The  action  (1)  is  of  a  type  very  common  in  chemistry.  It  is  more 
complex  even  than  double  decomposition  (p.  20),  and,  unlike  this,  its 
results  cannot  be  anticipated  by  guessing.  If  we  had  used  manganous 
oxide  MnO,  we  should  have  had  a  double  decomposition: 

MnO  +  2HC1  -*  H20  +  MnCl2,  (2) 

but  we  should  have  got  no  chlorine.  Perhaps  the  simplest  way  to 
describe  the  difference  between  these  two  actions  is  in  terms  of  the 
valence  of  the  manganese.  In  MnIVO2ir  the  element  is  quadrivalent. 
This  means  that  its  atomic  weight  professes  to  be  able  to  hold  four 
atomic  weights  of  a  univalent  element.  The  four  valences  of  oxygen 
(2O11)  can  do  the  same  thing.  In  equation  (1)  the  oxygen  fulfils  this 
promise  by  taking  4H1.  But  the  Mn^  can  hold  only  2C11,  per- 
manently, and  lets  the  other  2C11  go  free.  In  other  words,  the 
valence  of  the  atomic  weight  of  manganese  changes  in  the  course  of  the 
action.  In  equation  (2),  on  the  other  hand,  the  manganese  is  bivalent 
to  start  with  (MnnOn)  ,  and  is  able  to  retain  the  amount  of  chlorine 
(2C11)  equivalent  to  On.  Actions  like  that  of  manganese  dioxide  in 
(1)  are  classed  as  oxidations.  The  hydrogen  chloride,  or  rather  half 


220  INORGANIC   CHEMISTRY 

of  it,  is  oxidized.     A  graphic  mode  of  writing  may  make  this  remark 
clearer: 

Mniv  //  O  +  2HC1  ->  H20  +  Mn11  C12 


The  upper  half  is  a  double  decomposition,  the  lower  an  oxidation  by 
half  the  combined  oxygen  of  the  dioxide. 

In  practice,  instead  of  employing  aqueous  hydrochloric  acid  we 
frequently  use  the  materials  from  which  it  is  prepared,  namely,  com- 
mon salt  and  concentrated  sulphuric  acid  (p.  206),  along  with  the 
manganese  dioxide.  Under  those  circumstances,  the  action  appears 
more  complex,  but  is  simply  a  combination  of  the  two  chemical 
changes,  and  is  represented  by  the  equation: 

Mn02  +  2NaCl  +  3H2SO4  ->  2H2O  +  2NaHS04  +  MnSO4  +  2C1. 

Kinetic-  Molecular  View  of  these  Actions.  —  In  preparing 
chlorine  with  manganese  dioxide,  the  gas  is  produced  rather  slowly. 
The  situation  is  that  we  have  placed  together  manganese  dioxide  in 
a  granular  form  and  water  which  contains  hydrogen  chloride  in 
solution.  The  dioxide  is  very  insoluble  in  water,  and  consequently 
its  molecules,  which  must  dissolve  before  they  can  meet  the  acid, 
are  present  in  small  quantities:  Mn02  (solid)  <±  Mn02  (dslvd).  The 
finer  the  pulverization,  and  the  larger  the  amount  of  the  oxide,  the 
less  will  be  the  delay  from  this  cause.  On  the  other  hand,  the  acid 
contains  originally  only  about  one  molecule  of  hydrogen  chloride  for 
every  five  of  water  and,  as  the  former  is  used  up,  the  scarcity  of  the 
active  substance  becomes  greater. 

Again,  we  heat  the  mixture  on  a  water  bath  so  as  to  hasten  the 
process  (p.  93)  by  raising  the  temperature  to  about  90°.  When  we 
prepared  oxygen,  we  forced  the  temperature  up  with  a  naked  Bunsen 
flame  until,  at  about  300°,  a  sufficiently  rapid  stream  of  the  gas  was 
secured.  The  iron  and  sulphur  (p..  16)  we  raised  nearly  to  a  red  heat. 
Here  the  conditions  make  stronger  heating  impossible.  No  aqueous 
solution  of  hydrogen  chloride  can  be  raised  above  110°,  the  maximum 
boiling-point  (p.  211).  But  we  must  not  carry  the  heating  so  far  as 
110°,  because  even  below  this  point  the  concentrated  acid  gives  off 
gaseous  hydrogen  chloride  freely.  If  we  did,  we  should  contaminate 
our  chlorine  and  at  the  same  time  lose  a  part  of  one  of  the  ingredients 
on  which  the  action  depends.  Intelligent  chemical  work  always 
demands  a  careful  consideration  of  purely  physical  facts  of  this 
description. 


HYDROGEN  CHLORIDE  AND  CHLORINE  221 

On  the  other  hand,  when  potassium  permanganate  is  employed 

(p.  217),  the  chlorine  is  evolved  very  rapidly.  The  permanganate  is 
fairly  soluble  in  the  cold  (6.5  :  100  Aq  at  15°),  and  becomes  rapidly 
more  soluble  as  the  heat  of  the  reaction  raises  the  temperature  of  the 
liquid.  Then,  too,  the  permanganate  is  a  more  active  oxidizing  agent, 
at  equal  concentrations,  than  is  manganese  dioxide,  and  so  attacks  the 
hydrochloric  acid  more  rapidly. 

Physical  Properties,  —  Chlorine  differs  from  the  gases  we  have 
encountered  so  far  in  having  a  strong  greenish-yellow  tint  (Gk. 
xXwpos,  pale  green),  a  fact  which  gave  rise  to  its  name,  and  having  a 
powerful,  irritating  effect  upon  the  membranes  of  the  nose  and  throat. 

Density  (H  =  1),  35.79  Boiling-point  (liq.),  -33.6° 

Weight  of  22.4  L,  72.13  g.  Melting-point  (solid),  -102° 

Solubility  in  Aq  (20°),  215  vols.  in  100  Vap.  tension  (liq.)  0°,  3.66  atmos. 

Crit.  temp.,  +146°  Vap.  tension  (liq.)  20°,  6.62  atmos. 

Since  a  liter  of  air  weighs  1.293  g.,  chlorine  (wt.  1 1.,  3.22  g.)  is  two 
and  a  half  times  heavier.  In  solubility  it  stands  between  slightly 
soluble  gases,  like  oxygen  and  hydrogen,  and  those  which  are  ex- 
tremely soluble.  It  can  be  collected  over  hot  water  or  a  strong 
solution  of  salt. 

It  was  first  liquefied  by  Northmore  (1806).  The  critical  tempera- 
ture (p.  166)  is  exceptionally  high  (146°),  so  that  at  all  ordinary 
temperatures  the  gas  can  be  liquefied  by  compression  alone.  It  forms 
a  yellow  liquid  which,  contained  in  steel  cylinders,  is  now  an  article  of 
commerce.  On  being  cooled  below  —102°,  it  gives  a  pale-yellow 
solid. 

Chemical  Properties.  —  Chlorine  is  at  least  as  active  a  sub- 
stance as  is  oxygen.  It  presents  a"  more  varied  array  of  chemical 
properties  than  does  that  element  (see  below).  The  binary  com- 
pounds are  called  chlorides. 

Combines  with  Metals.  Powdered  antimony  (cold) ,  when  thrown 
into  chlorine,  unites  with  it  to  form  the  chloride  SbCls,  which  appears 
partly  as  vapor  and  partly  as  glowing  particles. 

Balanced:  Sb  +  3C1  ->  SbCl3. 

Copper,  in  the  condition  of  thin  leaf  commonly  used  for  gilding 
(Dutch-metal),  catches  fire  when  thrust  into  the  gas,  giving  a  fog 
of  solid  cupric  chloride  CuCl2.  Sodium  burns  brilliantly,  giving 
a  cloud  of  sodium  chloride.  The  union  of  a  metal  like  sodium  and 


222  INORGANIC  CHEMISTRY 

a  colored,  irritating  gas  to  give  a  mild  household  article,  like  common 
salt,  illustrates  the  extraordinary  nature  of  chemical  change.  All 
the  familiar  metals,  with  the  exceptions  of  gold  and  platinum,  com- 
bine readily  with  chlorine. 

When  metals  (like  copper  and  iron)  and  chlorine  are  first  thor- 
oughly freed  from  moisture,  combination  no  longer  occurs.  A  trace 
of  water  vapor  is  required  in  these,  as  it  is  in  many  other  chemical 
actions,  as  a  contact  agent.  Hence,  the  chlorine,  before  being 
compressed  into  steel  cylinders,  must  be  freed  entirely  from  water 
vapor  (see  Detinning). 

Combines  with  Hydrogen.  A  jet  of  hydrogen  burns  vigorously 
in  chlorine,  producing  hydrogen  chloride.  The  presence  of  this 
product  may  be  recognized  at  once,  because,  while  chlorine  in  contact 
with  moist  breath  gives  no  cloud,  hydrogen  chloride  (q.v.)'  produces 
a  dense  fog.  The  union  of  the  gases,  when  a  mixture  of  them  is  kept 
cold  and  in  the  dark,  is  too  slow  to  be  perceived.  On  exposure  to 
diffused  light,  however,  they  unite  slowly,  while  a  sudden  flash  of 
sunlight  or  the  burning  of  a  magnesium  ribbon  causes  instant  explo- 
sion. The  function  of  the  light  here  is  entirely  different  from  that  in 
the  decomposition  of  silver  chloride  (p.  19).  In  the  latter  case  light 
was  used  to  maintain  the  change,  which  comes  to  a  stop  whenever  the 
light  is  withdrawn :  the  action  was  endothermal  and  consumed  energy. 
The  union  of  hydrogen  and  chlorine  is  highly  exothermal,  and  a 
minimum  of  light  only  is  needed  to  start  it  (pp.  94-95).  The  action 
of  the  light  is  catalytic. 

Interacts  with  Compounds  Containing  Hydrogen.  When  a  lighted 
taper  is  plunged  into  chlorine  it  continues  to  burn,  but  a  dense  clou*d 
of  soot  (free  carbon)  rises  from  the  flame.  Blowing  the  breath  into 
the  jar  then  gives  the  fog  which  shows  the  presence  of  hydrogen 
chloride.  Thus  the  presence  of -hydrogen  and  carbon  in  the  wax  is 
proved.  We  learn,  also,  that  chlorine  has  little  tendency  to  combine 
with  carbon,  for  this  element  goes  free.  A  few  drops  of  warm  turpen- 
tine, poured  upon  a  strip  of  paper,  when  placed  in  chlorine  give  a 
violent  reaction  and  a  cloud  of  finely  divided  carbon  bursts  forth. 

Ci0H16  +  16C1  -»  16HC1  +  IOC. 

Elements  Displaced  by  Chlorine.  The  action  on  turpentine  is 
a  displacement  of  the  carbon  by  the  chlorine.  Of  the  same  nature 
is  the  action  of  chlorine  upon  potassium  iodide  KI,  dry  or  in  solution. 

KI  +  C1->KC1  +  I. 


HYDROGEN   CHLORIDE  AND  CHLORINE  223 

The  iodine,  when  moist,  is  deep  brown  in  color.  A  mere  trace  of 
chlorine,  liberating  a  trace  of  iodine,  gives  no  visible  effect.  But 
if  some  starch  is  present,  even  a  trace  of  free  iodine  yields  a  deep 
blue  color.  This  reaction  is  used  as  a  test  for  chlorine,  for  free 
iodine  from  any  source,  and  for  starch  (p.  5).  To  test  for  chlorine, 
strips  of  filter  paper,  dipped  in  starch  emulsion  (starch  boiled  with 
much  water  and  cooled)  to  which  a  few  drops  of  potassium  iodide 
have  been  added,  are  used.  Combined  iodine,  as  in  potassium 
iodide,  has  no  effect  upon  starch.  Combined  chlorine,  as  in  sodium 
chloride,  has  no  action  upon  the  prepared  strips  of  paper  —  free 
chlorine  is  required. 

Action  upon  Water.  We  have  seen  that  chlorine  seizes  the 
hydrogen  in  turpentine.  We  have  also  learned  that  it  combines 
with  the  hydrogen  in  steam,  reversing  Deacon's  process  to  the  extent 
of  20  per  cent.  It  also  acts  upon  cold  water,  when  dissolved  in  the 
latter,  although  in  a  similarly  incomplete  way.  The  substances 
formed  are  hydrochloric  acid  and  hypochlorous  acid  HC10  : 


With  half-saturated  chlorine-water  at  10°  -  -  that  is,  water  con- 
taining an  equal  volume  of  chlorine  gas  —  33  per  cent  of  the  chlorine 
is  changed  into  the  acids.  When  one-fifth  saturated  (10°),  52  per 
cent  is  changed  into  the  acids,  and  when  one-tenth  saturated,  73  per 
cent  (  Jakowkin)  .  The  percentages  become  greater  above  10°.  Thus, 
chlorine-water  (the  solution)  is  a  mixture  containing  dissolved 
chlorine  and  two  acids.  Hypochlorous  acid  (q.v.)  is  of  especial  interest 
because  it  is  a  very  active  substance,  with  powerful  oxidizing  qualities, 
and  bleaches  dyes  by  decomposing  them. 

The  action  comes  to  a  standstill  when  one-third  completed, 
because  the  two  acids  interact  to  reproduce  chlorine  and  water 
(read  the  equation  backwards).  The  action  is  reversible.  When 
the  solution  is  exposed  to  sunlight,  the  hypochlorous  acid  decom- 
poses and  oxygen  gas  is  liberated  and  escapes: 


Since  this  removes  the  hypochlorous  acid,  on  whose  interaction  with 
the  hydrogen  chloride  the  reverse  action  depends,  the  forward  action 
proceeds  under  continuous  illumination  gradually  to  completion. 
Hence  the  aqueous  solution  of  chlorine  must  be  kept  in  the  dark, 
since  otherwise,  after  a  time,  a  dilute  solution  of  hydrogen  chloride 
alone  remains. 


224  INORGANIC  CHEMISTRY 

The  reader  should  note  here  the  displacement  of  the  equilibrium, 
a  chemical  one  in  this  case,  in  consequence  of  the  annulment  of  one 
of  the  opposing  tendencies  (p.  170).  Through  the  destruction  of 
the  hypochlorous  acid,  one  of  the  tendencies,  namely,  that  repre- 
sented in  the  backward  action,  becomes  inoperative.  The  forward 
action  is  not  itself  assisted,  but  it  is  no  longer  impeded,  and  so  proceeds 
to  completion. 

The  so-called  bleaching  action  of  "chlorine"  is  almost  always  the  result  of  oxi- 
dation of  the  coloring  matter  by  hypochlorous  acid  (q.v.} .  Chlorine  and  the  dye  in 
the  cloth,  even  when  only  moderately  dry,  show  no  tendency  to  interaction.  This 
may  be  demonstrated  by  collecting  some  chlorine  in  a  stoppered  bottle  in  the 
bottom  of  which  a  little  concentrated  sulphuric  acid  stands.  A  piece  of  colored 
calico  may  be  attached  by  a  pin  to  a  cork  stuck  in  the  bottom  of  the  stopper  and 
so  suspended  in  the  gas.  After  twenty-four  hours  no  action  will  be  found  to  have 
occurred.  Yet  if  the  rag  is  first  moistened,  the  bleaching  is  almost  instantaneous. 
Hence  this  bleaching  action  is  treated  under  the  properties  of  hypochlorous  acid. 

So-called  "Nascent  Oxygen." — The  oxidizing  action  of  chlorine-water 
is  commonly  attributed  to  "nascent  oxygen"  —  oxygen  in  the  state  of  being 
born.  It  was  supposed  that  the  chlorine  preempted  the  hydrogen  of  the  water, 
without  actually  combining  with  it,  and  so  left  the  oxygen  in  a  distracted  condition 
in  which  it  was  more  active  than  free  oxygen!  But  all  this  speculation  occurred 
before  we  knew  that  there  was  hypochlorous  acid  in  chlorine-water.  A  pure 
solution  of  this  acid  performs  all  the  oxidations  that  chlorine-water  can  bring 
about.  So  that,  since  we  have  now  to  choose  between  a  substance  we  know  to 
be  present,  and  to  be  perfectly  capable  of  doing  the  work,  and  an  imaginary,  ill- 
defined  entity,  the  scientific  method  naturally  demands  that  we  attribute  the 
effect  to  the  substance.  In  this  connection,  the  correct  name  for  "nascent 
oxygen"  is,  therefore,  hypochlorous  acid. 

It  is,  in  any  case,  time  that  the  term,  and  the  idea  of  "nascent  oxygen" 
should  be  eliminated  from  the  science.  This  material  is  imaginary  —  it  has 
never  been  isolated  or  studied  quantitatively.  If  it  is  an  allotropic  form  of 
oxygen,  it  must  have  properties  and  a  degree  of  activity  that  can  be  defined 
quantitatively.  But  this  cannot  be  done,  because  it  has  not  always  the  same 
activity.  Thus,  if  all  oxidizing  agents  perform  their  oxidizing  by  means  of  "nas- 
cent oxygen,"  it  is  curious  that  hypochlorous  acid  oxidizes  hydrochloric  acid 
instantly  and  easily,  while  hydrogen  peroxide  does  not,  also  that  chloric  acid 
HClOs  oxidizes  hydrochloric  acid  rapidly,  while  perchloric  acid  HC1O4  does  not. 
(For  explanation  of  oxidation  by  oxidizing  agents,  see  pp.  314,  321.)  If  we  mean 
that,  when  the  free  elements  are  not  present  and  yet  compounds  containing  them 
interact,  we  must  assume  that  the  elements  are  in  the  nascent  condition,  then  we 
should  be  consistent,  and  explain  the  action  of  sulphuric  acid  on  sodium  chloride 
as  being  due  to  nascent  chlorine  and  nascent  hydrogen!  All  double  decomposi- 
tions would  demand  the  same  mode  of  explanation!  Finally,  since  every  oxida- 
tion is  accompanied  by  a  reduction,  when  we  assume  the  presence  of  nascent 


HYDROGEN  CHLORIDE  AND  CHLORINE  225 

oxygen,  to  be  consistent,  we  ought  to  assume  the  presence  of  nascent  hydrogen 
also:  each  such  case  involves  a  twin  birth.  The  conception  will  not  bear  careful 
examination. 

Action  by  Substitution.  When  actions  like  that  on  turpentine 
—  that  is  on  compounds  containing  carbon  and  hydrogen  —  are 
moderated  by  altering  the  conditions,  the  decomposition  is  not  so 
complete.  Using  a  lower  temperature  is  effective.  Thus,  if  methane 
CH4  (marsh-gas)  ,  the  chief  component  of  natural  gas,  is  mixed  with 
chlorine  and  exposed  to  sunlight,  a  slower  action  occurs,  of  which  the 
first  stage  consists  in  the  removal  of  one  unit  weight  of  hydrogen  and 
the  substitution  of  chlorine  for  it  according  to  the  following  equation: 

CH4  +  2C1  -i  CH3C1  +  HC1. 

The  process  may  continue  further  by  the  substitution*  of  chlorine 
for  the  units  of  hydrogen  one  by  one  until  carbon  tetrachloride  CCU 
is  finally  formed. 

The  action  on  water  (previous  section)  is  a  substitution. 

Combines  with  Non-metals.  Phosphorus  burns  in  chlorine  with 
a  rather  feeble  light,  producing  primarily  phosphorus  trichloride 
PC13,  a  liquid  (b.-p.  74°).  If  excess  of  chlorine  is  present,  then,  as 
the  trichloride  cools,  it  combines  to  form  the  solid  pentachloride 
PC15.  Sulphur,  when  heated,  unites  more  slowly,  giving  sulphur 
monochloride  S2C12,  a  liquid  used  in  vulcanizing  rubber.  Chlorine 
does  not  combine  directly  with  carbon,  nitrogen,  or  oxygen,  although 
compounds  with  those  elements  can  be  made  indirectly.  With  the 
helium  group  of  elements  (q.v.),  it  forms  no  compounds. 

Combines  with  Compounds.  Chlorine  unites  with  many  com- 
pounds. Thus,  one  of  the  oxides  of  carbon,  carbon  monoxide  CO, 
when  mixed  with  chlorine  and  exposed  to  sunlight  gives  drops  of  a 
volatile  liquid  (b.-p.  8.2°)  known  as  phosgene 


*  Substitution  resembles  displacement  (p.  18)  in  that  an  element  and  a 
compound  interact,  and  the  element  takes  the  place  of  one  unit  in  the  com- 
position of  the  latter.  In  the  above  action,  one  unit  of  chlorine  takes  the  place 
of  one-  unit  of  hydrogen.  But  the  latter  is  not  liberated;  it  combines  with  another 
unit  of  chlorine.  The  action  resembles  double  decomposition,  excepting  that 
one  of  the  substances  is  not  a  compound,  but  a  diatomic  element.  The  name 
used  is  intended  to  fix  the  attention  on  the  compound  and  on  the  fact  that  one 
unit  has  been  substituted  for  another  in  it.  This  conception  is  a  favorite  one 
in  the  chemistry  of  compounds  of  carbon. 


226 


INORGANIC   CHEMISTRY 


FIG.  90. 


When  chlorine-water  is  cooled  with  ice,  a  compound,  chlorine 
hydrate  C1,4H20  crystallizes  out.  Faraday  (1823)  placed  this  sub- 
stance in  the  closed  limb  of  a  A-tube,  sealed  the  open  end,  and  placed 
the  empty  limb  in  ice  and  water  (Fig.  90).  When  the  hydrate  was 
gently  warmed,  chlorine  gas  was  given  off 
and  was  liquefied  by  its  own  pressure  in  the 
cold  part  of  the  tube. 

Chemical  Relations  of  the  Element.* 

—  In  the  chlorides,  an  atomic  weight  of  chlo- 
rine is  equivalent  to  one  atomic  weight  of 
hydrogen  or  of  sodium.  The  element  is, 
therefore,  univalent  (p.  131).  It  never  shows  any  higher  valence  than 
this  save  in  its  oxygen  compounds  (see  Chap.  XXIII).  The  oxides 
of  chlorine  interact  with  water  to  give  acids,  and  the  element  is, 
therefore,  to  be  classed  as  a  non-metallic  element  (p.  150).  It  be- 
longs to  that  group  of  the  non-metallic  elements  called  the  halogens, 
as  a  consideration  of  some  others  of  its  relations  will  show  (see  Chap. 
XIV  and  Periodic  system). 

Uses  of  Chlorine.  —  Large  quantities  of  chlorine  are  manu- 
factured for  the  preparation  of  bleaching  materials  and  disinfecting 
agents.  In  disinfection,  the  minute  germs  of  disease  and  putrefaction 
are  acted  upon  by  the  hypochlorous  acid  formed  by  the  interaction 
of  chlorine  with  water,  and  instantly  their  life  is  destroyed. 

Chlorides.  —  The  chlorides  are  described  individually  under  the 
other  element  which  each  contains.  The  majority  of  the  chlorides 
of  the  metals  are  easily  soluble  in  water.  The  only  familiar  exceptions 
are  silver  chloride  AgCl,  mercurous  chloride  (calomel)  HgCl,  cuprous 
chloride  CuCl,.  aurous  chloride  (one  of  the  chlorides  of  gold)  AuCl, 
thallous  chloride  T1C1,  and  ordinary  lead  chloride  PbCl2.  The  last 
of  these  is  on  the  border  line  as  regards  solubility.  An  appreciable 

*  In  accordance  with  the  distinction  that  must  be  drawn  (p.  23)  between  the 
element  as  a  variety  of  matter  in  combination,  and  the  elementary  substance  or  free 
form  of  the  element,  and  to  avoid  a  common  source  of  confusion,  we  shall  always 
give  only  the  behavior  of  the  elementary  substance  under  the  title  chemical  proper- 
ties. The  characteristics  which  distinguish  the  compounds  of  the  element,  as  a 
class  from,  or  relate  them  as  a  class  to  the  compounds  of  other  elements  will  then 
appear  in  a  separate  section  under  the  title  "chemical  relations"  (see  Chap.  XIV, 
first  section  and  Chap.  XXII,  Periodic  system,  fifth  section). 


HYDROGEN   CHLORIDE  AND  CHLORINE 


227 


amount  dissolves  in  cold  water,  and  a  considerable  amount  in  boiling 
water.     For  the  various  modes  of  preparing  chlorides,  see  p.  213. 

Composition  of  Hydrogen  Chloride.  —  Being  now  familiar 
with  both  hydrogen  and  chlorine,  we  may  take  up  the  proportions 
by  weight  in  which  they  combine,  and  also  the  proportions  by  volume 
in  which  the  constituents  unite,  and  the  relation  of  this  to  the  volume 
of  the  resulting  hydrogen  chloride. 

The  proportion  of  hydrogen  to  chlorine  by  weight  in  this  com- 
pound is  1  :  35.18.  Taking  the  atomic  weight  of  hydrogen  1.008,  so 
as  to  harmonize  that  of  chlorine  with  O  =  16,  the  standard  scale  for 
atomic  weights,  the  ratio  becomes  1.008  :  35.46. 

The  proportion  by  volume  in  which  the  constituents  unite,  and 
the  relation  of  this  to  the  volume  of  the  resulting  hydrogen  chloride, 
may  easily  be  shown  in  several  ways.  The  decomposition  of 
the  solution  of  hydrogen  chloride  in  water 
by  means  of  the  electric  current  proves 
that  the  gases  are  liberated  in  equal 
volumes. 

The  apparatus  in  Fig.  44  (p.  120)  can- 
not be  used  to  show  this,  because,  under 
the  increasing  pressure  due  to  the  dis- 
placement of  the  liquid  into  the  higher 
bulb,  the  chlorine  becomes  more  and  more 
soluble,  and  its  volume  therefore  falls  pro- 
gressively more  and  more  below  what  it 
should  be. 

Brownlee's  apparatus  for  demonstrat- 
ing this  is  shown  in  Fig.  91.  The  central 
part  is  the  same  as  in  Fig.  44,  but,  when 
the  three-way  stopcock  is  closed,  the  gases 
go  to  right  and  left,  and  displace  the  liquid 
in  two  outside  tubes  (water  in  one  and 
chlorine- water  in  the  other).  The  equal 
rate  at  which  this  takes  place  on  both  sides 
proves  that  the  gases  are  generated  in  equal 
volumes. 

In  order  to  ascertain  the  relation  between  the  volumes  of  the 
constituents  and  that  of  the  product,  we  may  unite  the  gases  and  find 
out  whether  any  change  in  volume  occurs.  A  tube  with  thick  walls 
(Fig.  92)  is  filled  with  the  mixed  gases  obtained  by  electrolysis.  By 


FIG.  91. 


228  INORGANIC   CHEMISTRY 

dipping  one  end  of  the  tube  under  mercury  and  opening  the  lower 
stopcock,  it  is  seen  that  no  gas  leaves  and  no  mercury  enters.     After 
the  mixture  has  been  exploded,  by  the  light  from  burning  magnesium, 
.n         the  same  test  is  repeated  with  the  same  result.     The  pres- 
sure has  therefore  remained  equal  to  that  of  the  atmosphere, 
f^       Hence  there  has  been  no  change  in  volume  as  the  result  of 
the  union.     It  appears  therefore,  that: 

1  vol.  hydrogen  -f-  1  vol.  chlorine  -*  2  vols.  hydrogen  chloride, 
a  result  in  harmony  with  Gay-Lussac's  law  (p.  157). 

Another  way  of  demonstrating  the  equality  in  the  volumes  of  the 
hydrogen  and  chlorine  is  to  fill  a  wide  tube,  closed  at  each  end  by  a 
stopcock,  with  the  mixed  gases  arising  from  electrolysis  of  hydrochloric 
acid.     When  the  air  has  been  entirely  displaced,  the  stopcocks  are 
closed.     The  gases  which  the  tube  then  contains  are  present  very 
nearly  in  the  proportions  in  which  they  are  liberated  from  the  decom- 
'IG>      '      position  of  the  substance.    By  introducing  a  small  amount  of  potassium 
iodide  solution,  the  chlorine  is  removed  (KI  +  Cl   --»  KC1  +1).     It  forms  potas- 
sium chloride,  which  remains  dissolved  in  the  water,  and  free  iodine,  which  dis- 
solves in  the  excess  of  potassium  iodide  solution.     Neither  product  is  gaseous 
under  the  circumstances,  so  that  the  volume  of  the  mixed  gases  diminishes  by 
the  amount  of  chlorine  removed.     If  the  stopcock  is  now  opened  under  water,  the 
latter  enters  and  fills  half  the  length  of  the  tube.     The  remaining  gas  is  easily 
shown  to  be  hydrogen. 

To  show  that  the  volume  of  the  hydrogen  chloride  is  twice  that  of  either 
constituent  in  the  free  condition,  an  alternative  method  is  likewise  available. 
We  may  completely  fill  a  long  test-tube  with  hydrogen  chloride,  introduce  into  it 
quickly  some  sodium  dissolved  in  mercury,  and,  after  agitation,  open  the  closed 
tube  under  mercury.  The  sodium  gives  sodium  chloride  and  free  hydrogen,  and 
it  is  found  that  the  mercury  enters  so  as  to  fill  one-half  of  the  tube.  Since  from 
this  experiment  we  learn  that  the  hydrogen  occupies  half  the  space  of  the 
hydrogen  chloride,  and  from  the  previous  experiment  we  know  that  the  volume 
of  hydrogen  is  equal  to  that  of  the  chlorine,  we  conclude  that  two  volumes  of 
mixed  hydrogen  and  chlorine  would  give  two  volumes  of  hydrogen  chloride. 

Classification  of  Chemical  Interactions  and  Exercises 
Thereon.  —  So  far  we  have  defined  ten  more  or  less  distinct  kinds 
of  chemical  change:  Combination  (p.  12),  decomposition  (p.  17), 
dissociation  (p.  148),  displacement  (p.  18),  substitution  (p.  225), 
double  decomposition  (p.  20),  hydrolysis  (p.  210),  oxidation  (pp. 
91,  217,  218),  reduction  (pp.  92,  127),  and  electrolysis  (p.  121).  In  one 
or  two  of  these  classes  all  the  actions  are  reversible,  in  others  some  are 
reversible  and  some  are  not.  Illustrations  of  every  one  of  these  will  be 


HYDROGEN   CHLORIDE  AND  CHLORINE  229 

found  in  the  present  chapter.  The  classes  are  not  mutually  exclusive. 
Some  actions  belong  to  one  class  or  another  according  to  our  point 
of  view  at  the  moment.  The  ability  readily  to  classify  each  phe- 
nomenon, as  it  comes  up,  requires  precisely  that  grasp  of  the  frame- 
work of  the  science  which  the  reader  must  seek  speedily  to  attain. 
For  example,  let  him  classify  the  following  actions:  1.  Action  of 
heat  on  chloroplatinic  acid;  2.  of  potassium  on  water;  3.  of  heat  on 
potassium  chlorate;  4.  of  chlorine  on  metals;  5.  of  chlorine  on  tur- 
pentine; 6.  of  chlorine  on  potassium  iodide;  7.  of  chlorine  on  meth- 
ane; 8.  of  carbon  monoxide  and  chlorine;  9.  of  sunlight  on  hypo- 
chlorous  acid;  10.  of  sulphuric  acid  on  salt;  11.  of  zinc  oxide  and 
hydrochloric  acid;  12.  of  zinc  on  hydrochloric  acid. 

13.  Expand  the  explanation  of  the  tendency  of  hydrogen  chloride 
to  fume  in  moist  air  (p.  211). 

14.  Explain  the  interaction  of  steam  and  iron  (p.  128)  on  mechani- 
cal principles  similar  to  those  used  in  describing  how  hydrogen  chlo- 
ride is  formed  from  salt  and  sulphuric  acid  (p.  208). 

15.  In  the  interactions  of  potassium  permanganate  and  of  man- 
ganese dioxide,  respectively,  with  hydrochloric  acid,  what  fractions 
of  the  whole  chlorine  are  liberated?     What  are  the   commercial 
advantages  of  the  use  of  salt  and  sulphuric  acid  with  the  manganese 
dioxide? 

16.  What  are  the  relative  volumes  of  the  gaseous  interacting 
substances  and  products  in  the  following  reactions:    (a)  turpentine 
vapor  and   chlorine;     (6)   methane  and  chlorine;     (c)    phosphorus 
vapor  and  chlorine;   (d)  carbon  monoxide  and  chlorine. 

17.  In  view  of  the  explanations  given,  can  you  define  the  general 
nature  of  the  " other  substances"  (p.  218)  which  may  be  used  to 
oxidize  hydrochloric  acid? 

SUMMARY  OF  PRINCIPLES 

It  may  be  useful  at  this  point  partially  to  summarize  the  principles  (general 
facts)  of  chemistry  so  far  as  they  have  been  developed  in  the  preceding  chapters. 
These  principles  are  given  under  fourteen  heads  below.  They  are  stated  as  far 
as  possible  strictly  in  terms  of  facts,  since  hypotheses  are  not  integral  parts  of 
chemistry,  but  are  scaffolding  temporarily  employed  to  facilitate  the  erection  of 
the  structure  of  the  science.  In  a  later  chapter  (Chap.  XV),  some  other  impor- 
tant principles  will  be  summarized  in  like  manner.  To  secure  more  strictly  logical 
arrangement  than  has  seemed  advisable  in  the  text,  two  conceptions  which  have 
already  been  dealt  with  are  held  over  to  the  second  half  of  the  summary,  namely, 
17  (valence)  and  21  (chemical  relations  of  elements).  The  reader  should  give  care- 
ful thought  to  the  various  points,  many  of  which;  in  a  backward  view,  will  be  found 


230  INORGANIC  CHEMISTRY 

to  have  become  susceptible  of  improved  statement.  We  begin  the  series  with  the 
most  fundamental  fact  of  all,  —  the  one  without  which  no  chemical  work  would  be 
possible: 

1.  Each  substance  has  its  own  set  of  specific  physical  properties.     By  means 
of  these  it  is  recognized  and,  when  necessary,  separated  from  other  substances 
(p.  3). 

2.  Substances  are  either  simple  (elementary),  containing  only  one  kind  of 
matter,  or  compound,  containing  more  than  one  kind  of  matter  (p.  23). 

3.  In  all  chemical  actions  (excepting  "internal  rearrangements"),  changes 
in  the  material  composition  of  bodies  occur  (p.  21). 

4.  In  chemical  actions  there  is  no  change  in  the  total  mass  of  the  system 
(p.  52). 

5.  Each  substance  has  a  definite  material  composition  by  weight  (p.  54). 

6.  The  proportions  by  weight  in  which  all  chemical  combinations  take  place 
can  be  expressed  in  terms  of  small  integral  multiples  of  fixed  numbers,  which  may 
be  called  combining  weights,  one  for  each  element.     That  weight  of  each  element 
which  combines  with  8  parts  of  oxygen  is  called  the  equivalent  weight  and  has  the 
properties  of  a  combining  weight  (pp.  61,  63). 

7.  The  proportions  by  volume  in  which  all  chemical  interactions  involving 
substances  in  the  gaseous  condition  take  place  are  expressible  by  small  integers 
(p.  157). 

8.  From  6  and  7  it  follows  that  the  equivalent  weights  of  all  substances  oc- 
cupy, in  the  gaseous  condition  and  at  the  same  temperature  and  pressure,  volumes 
which  are  either  equal  or  stand  to  one  another  in  the  ratio  of  small  integers  (p.  158). 

9.  In  every  chemical  phenomenon  a  transformation  of  energy  occurs.     This 
results  in  a  redistribution  of  the  internal  energy  in  the  substances  concerned,  and 
also  in  an  increase  or  a  decrease  in  the  free  internal  energy  in  the  system  (p.  35). 

10.  In  chemical  phenomena  there  is  no  actual  loss  or  gain,  but  only  transfor- 
mation of  energy  (p.  32). 

11.  Interactions  which  proceed  spontaneously  are  in  general  those  in  which 
the  free  energy  is  transformed  into  some  other  variety  or  varieties  of  energy 
(p.  35). 

12.  Each  substance  has  its  own  set  of  chemical  properties,  such  as: 

(a)  Affinity:  the  given  substance  can  or  can  not  interact  with  such  and  such 
elementary  and  compound  substances. 

(6)  Relative  activity  of  the  systems  in  (a).  This  is  measured  quantitatively 
by:  (a)  Relative  speed  under  like  conditions  (see  13,  14,  18);  (/?)  Relative  heat 
developed,  when  actions  compared  can  be  carried  out  so  that  all  conditions  are 
alike;  (7)  Relative  E.M.F.  of  cell  when  the  action  is  so  arranged  as  to  give 
electricity  (pp.  128,  98,  37). 

13.  The  speed  of  every  interaction  is  increased  by  raising  the  temperature 
(p.  93). 

14.  The  speed  of  interactions  is  increased  or  decreased  by  catalytic'  agents, 
each  of  which  is  individual  in  the  kind  and  amount  of  its  effect  (p.  97). 


CHAPTER  XII 
MOLECULAR  WEIGHTS  AND  ATOMIC   WEIGHTS 

AVOGADRO'S  law  (p.  163)  has  proved  to  be  by  far  the  most  sug- 
gestive and  fruitful  of  all  the  conceptions  developed  from  the  kinetic- 
molecular  theory.  We  are  now  in  a  position  to  discuss  several  of 
its  most  important  applications.  These  concern  more  particularly 
the  measurement  of  the  relative  weights  of  the  molecules  of  different 
gaseous  substances,  and  the  determination  of  the  most  convenient 
magnitudes  for  the  chemical  unit  weights  (atomic  weights;  cf.  p.  63). 

Meaning  of  Avogadro's  Hypothesis.  —  First,  we  must  under- 
stand clearly  what  is  implied  in  the  statement  that :  In  equal  volumes 
of  all  gases,  at  the  same  temperature  and  pressure,  there  are  equal 
numbers  of  molecules.  It  means  that,  for  instance,  at  100°  and  760 
mm.,  in  specimens  of  all  gases,  the  average  spacing  of  the  molecules 
is  identical.  This  condition  is  independent  of  the  nature  of  the  gas 
—  for  example,  whether  it  is  a  simple  or  a  compound  substance,  like 
oxygen  and  carbon  dioxide,  respectively,  or  a  mixture,  like  air.  It 
means  that  when,  at  some  fixed  temperature,  we  fill  the  same  vessel 
with  a  number  of  different  gases  or  gaseous  mixtures  successively, 
the  number  of  molecules  that  it  will  hold  at  a  pressure,  say,  of  one 
atmosphere  will  always  be  the  same.  If  we  take  care  to  keep  tempera- 
ture and  pressure  the  same,  the  equality  in  the  number  of  molecules 
that  will  enter  the  jar  will  take  care  of  itself  automatically.  In 
what  follows,  to  avoid  continual  repetition,  standard  conditions  (0° 
and  760  mm.)  are  assumed  unless  other  conditions  are  specifically 
mentioned. 

This  statement  would  be  strictly  true  only  in  the  case  of  gases,  if  such  existed, 
which  behaved  in  ideal  accord  with  the  laws  of  Boyle  and  Charles.  Since,  how- 
ever, in  all  gases,  with  the  exception  of  hydrogen,  a  certain  tendency  to  cohesion 
between  molecules  is  distinctly  noticeable  and  its  amount  varies  from  gas  to  gas, 
the  density  with  which  the  molecules  are  packed  is  not  precisely  the  same  in  any 
two  of  them  (p.  164).  Hence,  Avogadro's  law  is  not  perfectly  realized  in  any 
known  gases.  In  the  case  of  hydrogen,  for  example,  a  very  slight  divergence 
from  the  behavior  of  an  ideal  gas  exists,  and  in  the  case  of  chlorine  it  amounts 

231 


232  INORGANIC  CHEMISTRY 

to  about  1^  per  cent,  and  is  quite  conspicuous.  This  slight  irregularity  in  the 
packing  of  the  molecules,  however,  does  not  interfere  with  the  application  of 
this  law  in  chemistry. 

Two  Kinds  of  Laws  in  Science.  —  It  will  have  been  observed 
that  the  laws  of  science  may  be  divided  into  two  kinds.  Some,  like 
those  of  conservation  of  mass  and  of  definite  proportions,  express  the 
facts  with  perfect  exactness.  The  divergence  between  our  experi- 
mental data  and  these  laws  we  find  to  become  smaller  and  smaller 
the  more  carefully  our  experiments  are  made.  The  difference  be- 
tween our  best  determinations  and  the  ideal  described  by  the  law, 
is  always  less  than  the  known  errors  of  observation.  There  is,  how- 
ever, a  second  class  of  laws  in  which  the  opposite  is  the  case.  The 
more  carefully  our  measurements  are  made,  the  more  clearly  is  it 
recognized  that  the  second  kind  of  law  does  not  state  the  fact  with 
exactness  for  any  single  example.  The  known  errors  of  measure- 
ment are  in  these  cases  less  than  the  discrepancy  between  the  ob- 
served fact  and  the  ideal.  The  laws  of  Boyle,  of  Charles,  and  of 
Gay-Lussac  are  examples  of  this  class.  The  value  of  such  laws  is 
not  impaired  by  the  fact  that  actual  substances  are  in  no  case  accu- 
rately described  by  them.  The  law  gives  us  a  norm  of  behavior  to 
which  most  of  the  substances  conform  with  a  fair  degree  of  closeness. 
Such  a  law  resembles  a  limit  in  mathematics,  toward  which  some 
expression  tends  to  converge  although  it  does  not  actually  reach  it. 
The  first  kind  of  law  represents  the  actual  behavior  of  materials,  the 
second  kind  of  law  an  average  behavior  to  which  no  one  material 
adheres  with  perfect  exactness. 

Now  Avogadro's  law  belongs  to  the  second  class.  No  known 
gases  conform  to  it  with  perfect  strictness. 

MOLECULAR  WEIGHTS 

The  Relative  Weights  of  the  Molecules.  —  According  to  Avo- 
gadro's hypothesis,  vessels  of  equal  size  filled  with  different  gases 
contain  equal  numbers  of  gaseous  molecules.  Now  equal  volumes  of 
different  gases  differ  very  markedly  in  weight,  or,  in  other  words, 
the  densities  of  various  known  gases  cover  a  wide  range  of  values. 
Thus,  hydrogen  is  the  lightest  of  all,  chlorine  is  more  than  thirty- 
five  times,  mercuric  chloride  (corrosive  sublimate)  vapor  over  one 
hundred  and  thirty-four  times  as  heavy.  Since  these  different  weights 
of  equal  volumes  represent  the  weights  of  equal  numbers  of  mole- 
cules, the  difference  must  be  due  to  the  differing  weights  of  the 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS 


233 


molecules  themselves.      The  densities   of  gases,   therefore,  may  be 
taken  as  measures  of  the  relative  weights  of  their  individual  molecules. 

The  extreme  significance  of  this  inference  in  chemistry  will  appear- 
as  we  elaborate  upon  it. 

The  various  scales  on  which  the  densities  *  of  gases  may  be  cal- 
culated, such  as  the  weights  of  one  liter  of  each  gas,  or  the  weights 
of  volumes  equal  to  that  of  one  gram  of  air,  are  illustrated  in  the 
first  two  columns  of  the  following  table: 


Weightf  of 
One  Liter,  0° 
and   760  mm. 

Density 
Air  =  1. 

Molecular 
Weight, 
Ox.  =  32. 

Hydrogen  .  •  

0.090 

0.0696 

2.016 

Oxygen  

1.439 

1.105 

32.00 

Chlorine 

3  166 

2  449 

70  92 

Hydrogen  chloride  
Carbon  dioxide  ... 

1.628 
1  965 

1.259 
1.520 

36.468 
44.00 

Water  
Mercury  
Mercuric  chloride  
Air  

0.8045 
8.957 
12.121 
1.293 

0.622 
6.908 
9.354 
1.00 

18.016 
200.6 
271.52 
28.955 

The  values  for  water  (b.-p.  100°),  mercury  (b.-p.  357°),  and  mer- 
curic chloride  (b.-p.  305°)  are  measured  at  high  temperatures  and 
reduced  by  rule  (p.  110)  to  0°  and  760  mm.  All  the  numbers  in  the 
first  two  columns,  as  they  stand,  are  purely  physical  in  derivation. 
Those  in  the  second  column  are  obtained  from  those  in  the  first  by 
using  the  proportion: 

1.293  (wt.  1  1.  air)  :  1.00  (air  =  1)  :  :  wt.  of  1  1.  any  gas  :  x 

(where  x  =  dens,  of  that  gas  on  scale  air  =1). 

The  last  column  will  be  explained  presently.  Since  the  numbers 
in  the  first  column  apply  to  equal  volumes  (1  1.),  and  those  in  the 
second  stand  in  constant  ratios  to  them,  the  weights  in  the  second 
column  represent  equal  volumes  also.  In  the  second,  the  volume 
is  Trshs-s  1.  The  values  in  either  one  of  the  columns  represent  the 


*  To  speak  strictly,  the  density  of  a  gas  is  the  weight  of  1  c.c.  at  0°  and  760  mm. 
Its  value  for  each  gas  is,  therefore,  obtained  by  dividing  the  numbers  in  the  first 
column  by  1000.  The  numbers  thus  obtained  are,  however,  inconveniently  small, 
and,  besides,  the  beginner  usually  measures  the  weight  of  a  liter  of  several  gases  in 
the  laboratory,  and  so  is  more  accustomed  to  the  unit  employed  above. 

t  These  are  not  the  observed  weights.  The  values  have  been  corrected  for 
the  two  deviations  from  the  laws  of  gases  (pp.  164-165). 


234 


INORGANIC  CHEMISTRY 


relative  weights  of  the  molecules  of  the  various  substances  (see 
Exercise  1  in  this  chapter). 

In  order  to  avoid  the  creation  of  unnecessary  confusion  in  the  mind  of  the 
beginner,  the  weights  of  one  liter  of  gas  in  the  above  table  are,  with  the  exception 
of  that  of  oxygen,  all  corrected  for  the  deviations  from  the  laws  of  Boyle  and 
Charles  (p.  164).  This  enables  us  to  show  the  process  by  which  the  molecular 
weights  are  derived  from  the  weights  of  one  liter  without  the  exhibition  of 
arithmetical  discrepancies  which  might  obscure  the  principle  being  explained. 
The  weights  of  one  liter  of  the  various  gases,  as  we  have  given  them,  are  based 
on  the  assumption  that  the  molecules  are  always  packed  uniformly  in  accordance 
with  Avogadro's  hypothesis.  The  actual,  measured  values  are  in  most  cases 
somewhat  different  from  these,  and  we  attribute  the  divergencies  to  the  varying 
degrees  of  cohesion  between  the  molecules  of  different  substances.  Even  the 
weight  of  one  liter  of  the  same  gas,  after  reduction  to  0°  and  760  mm.,  is  found  to 
vary  with  the  temperature  and  pressure  at  which  it  was  examined.  This  is  but 
natural,  since  changes  in  these  conditions  alter  the  effects  of  cohesion.  The 
following  table  gives  the  actual  weights  of  one  liter  of  the  same  gases,  with  a  few 
additional  ones,  and  a  comparison  will  show  the  extent  of  the  divergencies.  The 
most  interesting  case  perhaps  is  that  of  oxygen  and  hydrogen.  The  chemical 
combining  weights  of  these  substances  are  in  the  ratio  of  15.88  :  1.00,  while  a 
slight  excess  of  cohesion  in  oxygen  gives  the  ratio  of  their  densities  the  value 
15.90  :  1.00.  These  numbers  are  in  both  cases  based  upon  Morley's  results. 


Weight  of 
One  Liter,  0° 
and  760  mm. 

Density, 
Air  =  1. 

Observed  Molec- 
ular Weight, 
Ox.  =  32. 

Adjusted 
Molecular 
Weight. 

Hydrogen 

0  08987 

0  0695 

2  01 

2  016 

Oxygen 

1  429 

1  105 

32.00 

32.00 

Nitrogen    •  ';""  .  •. 

1  2507 

0  967 

28  02 

28  02 

Chlorine 

3  220 

2  490 

72  01 

70  92 

Hydrogen  chloride  .    .    . 
Carbon  dioxide      .    .^  .   .' 
Hydrogen  sulphide  .  ,-.  j  . 
Ammonia     ...*... 

1.6398 
1.9768 
1.537 
0.7708 

1.269 
1.529 
1.189 
0  597 

36.72 
44.28 
34.43 
17.27 

36.468 
44.00 
34.976 
17.034 

Sulphur  dioxide     .... 
Water   

2.9266 
0.8322 

2.264 
0.643 

65.56 
18.63 

64.06 
18.016 

Mercury      

8.87 

6.86 

198.4 

200.6 

Air 

1  293 

1.00 

28  955 

[Mixture] 

Molecular  Weights.  —  The  foregoing  section  shows  that,  pro- 
vided a  substance  is  a  gas  or  can  be  volatilized,  the  relative  weight 
of  its  molecules,  compared  with  those  of  other  volatile  substances, 
can  be  ascertained.  To  save  words,  the  relative  weight  of  the  mole- 
cule of  a  substance  is  called  the  molecular  weight  of  the  substance. 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          235 

Since  the  absolute  weights  of  molecules  cannot  be  accurately  de- 
termined, the  next  question  which  arises  is  as  to  the  choice  of  an 
appropriate  unit,  and  therefore  an  appropriate  scale  for  these  relative 
weights,  or  molecular  weights.  Now  the  numbers  already  given,  in 
the  first  two  columns  of  the  table,  are  purely  physical  data  and,  as 
they  stand,  lack  direct  relation  to  chemical  facts.  A  set  of  chemical 
numbers  is  required  for  chemical  purposes.  We,  therefore,  proceed 
next  to  show  how  the  required  relation  between  the  relative  weights 
of  molecules  and  chemical  facts  can  be  established. 

Chemistry  deals  with  chemical  combination,  and  most  substances 
are  compounds.  If  we  fix  our  attention,  then,  first,  on  compound 
substances  and  their  molecules,  we  perceive  at  once  that  the  mole- 
cules of  a  compound  substance  must  contain  two  or  more  elements 
in  definite  proportions  by  weight.  Thus  the  molecule  of  the  com- 
pound will  contain  one  or  more  atoms  of  each  of  the  component  elements. 
This  conception  at  once  suggests  two  ideas.  In  the  first  place,  since 
we  can  now  determine  the  relative  weights  of  molecules,  we  should 
also  somehow  be  able,  with  the  help  of  the  combining  proportions, 
to  determine  the  relative  weights  of  the  atoms  of  the  elements.  If 
these  relative  weights  of  atoms  are  properly  determined,  then  the 
weights  of  the  atoms,  when  added  together,  should  give  the  weight 
of  the  molecule.  Thus  —  and  this  is  the  second  idea,  and  the  one 
of  most  immediate  use  to  us  —  the  scale  for  relative  weights  of 
molecules  must  be  chosen  with  reference  to  the  scale  for  relative 
weights  of  the  atoms,  so  that  the  former  weights  may  always  include 
the  latter.  That  is  to  say,  the  molecular  weights  must  be  based 
upon  the  equivalent  weights  (p.  63).  This  means  that  the  scale 
must  be  such  that  no  molecule  of  a  compound  of  hydrogen  shall 
receive  a  value  so  small  that  the  proportion  of  hydrogen  in  it  is  less 
than  1.008.  Furthermore,  since  the  atomic  weight  of  oxygen  is  the 
standard  for  combining  proportions,  it  is  desirable  to  use  this  sub- 
stance as  basis  of  the  scale  of  molecular  weights.  Now,  as  we  shall 
find  (see  p.  240),  it  turns  out  that,  to  avoid  obtaining  proportions 
of  hydrogen  less  than  unity,  we  are  compelled  to  take  the  scale  32 
for  the  molecular  weight  of  oxygen.  Our  chemical  scale  of  densities 
is  therefore  calculated  to  the  scale,  density  of  oxygen  =  32.  This, 
then,  is  the  answer  to  the  problem  with  which  the  section  opened. 
The  third  column  of  the  table  (p.  233)  shows  the  results  of  re- 
calculating the  densities  to  this  chemical  scale.  The  pronortion 
used  is: 

Density  of  Ox.  :  Density  of  Substance  : :  32  :  x. 


236  INORGANIC  CHEMISTRY 

Thus,  if  we  take  the  densities  from  the  first  column,  the  value  for 
water  is  found  by  the  proportion,  1.429  :  (X8045  :  :  32  :  x  (=  18.016). 
That  is,  we  multiply  the  weight  of  a  liter  of  the  gas  by  32/1.429  to 
get  the  molecular  weight. 

Since  the  gram  is  the  unit  of  weight,  32  g.  of  oxygen,  or  18.016  g. 
of  water  is  called  the  gram-molecular  weight  of  the  substance.  It 
will  be  noted  that  32  g.  is  not  the  weight  of  a  molecule  of  oxygen. 
It  is  the  weight  of  a  very  large  number  of  molecules  of  oxygen  (see 
p.  238).  But,  whatever  that  number  of  molecules  of  oxygen  is, 
18.016  g.  of  water  contains  the  same  number  of  molecules,  and  the 
other  weights  in  the  same  column  are  weights  of  numbers  of  mole- 
cules equal  to  these.  The  term  gram-molecular  weight  being  some- 
what ponderous,  we  abbreviate  it  to  molar  weight,  and  still  further 
to  mole.  Thus,  a  mole  of  chlorine  is  70.92  g.  of  the  element,  and 
a  mole  of  hydrogen  chloride  is  36.468  g.  of  the  compound. 

When  the  above  method  of  calculating  the  molecular  weight  from  the  weight 
of  one  liter  of  a  substance  is  applied  to  the  actual  experimental  values,  the  result- 
ing molecular  weights  necessarily  diverge  somewhat  from  the  ideal  ones  which  we 
have  given.  Thus,  since  a  liter  of  hydrogen  chloride  actually  weighs  1.6398,  this 
number  when  multiplied  by  32  and  divided  by  1.429  gives  the  value  36.72.  This 
is  because  the  tendency  to  cohesion  causes  1  liter  of  this  gas  to  contain  more 
molecules  than  does  1  liter  of  oxygen.  The  mole  is  always  the  adjusted  molecular 
weight  and  not  the  observed  one. 

The    Gram-Molecular     (Molar)     Volume:      G.M.V.  —  The 

weights  in  the  last  column  of  the  table  (p.  233)  must  represent  equal 
volumes  of  the  different  gases.  This  follows  from  the  fact  that  they 
are  derived  from  the  values  in  the  first  column  by  multiplying  by  a 
constant  ratio  (32/1.429),  and  the  volume  in  the  first  column  is  always 
1  liter.  The  actual  dimension  of  this  volume  is  evidently  32/1.429 
liters,  which  is  almost  exactly  22.39,  or  in  round  numbers  22.4  liters. 
This  volume  at  0°  and  760  mm.  holds  32  g.  of  oxygen,  70.92  g.  of 
chlorine,  44.00  -g.  of  carbon  dioxide,  or,  in  fact,  the  molar  weight  of 
any  gaseous  substance.  It  is  called,  therefore,  the  gram-molecular 
volume  (G.M.V.)  or  the  molar  volume.  It  may  be  defined  as  that 
volume  which  contains  one  mole  (gram-molecular  weight)  of  any  gas 
at  0°and  760  mm.  At  other  temperatures  and  pressures  the  G.M.V. 
has  correspondingly  different  volumes. 

The  G.M.V.  gives  us  a  concrete  conception  of  a  molar  weight. 
This  volume  is  represented  by  a  cubical  wooden  box  (Fig.  93)  28.19 
cm.  (or  about  11.1  inches)  high.  Like  any  other  volume,  it  holds 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          237 

identical  numbers  of  molecules  of  different  gases.*  Its  capacity  at 
0°  and  760  mm.  is  the  number  of  molecules  in  32  g.  of  oxygen.  Hence, 
in  terms  of  molecules,  the  weight  of  any  gas  which  fills  it  bears  to 
32  g.  the  same  ratio  as  the  weight  of  a  mole- 
cule of  that  gas  to  the  weight  of  a  molecule 
of  oxygen. 


G.M.V. 

22.4  LITERS 


Measurement  of  Molar  Weights  (Moles)  . 

-  We  may  now  state  the  method  of  finding  the 
molar  (gram-molecular)  weight  of  a  substance 
thus:  Weigh  a  known  volume  of  the  substance, 
at  any  temperature  and  pressure  at  which  it  is  FlG 

gaseous,  reduce  this  volume  by  rule  to  0°  and 

760  mm.,  and  calculate  by  proportion  the  weight  of  22.4  liters  (see 
Exercises  1,  2,  3,  4). 

That  quantity  of  each  substance,  simple  or  compound,  which  at 
0°  and  760  mm.  would  fill  the  G.M.V.  cube  is  the  unit  quantity  of  the 
substance  for  all  theoretical  purposes  in  chemistry.  It  represents  the 
relative  weight  of  the  molecules  of  the  substance.  We  shall  employ 
it  at  once  for  the  purpose  of  determining  the  relative,  weights  of  atoms, 
or  atomic  weights. 

It  is  evident  that  the  chemical  molecular  weights,  adjusted  so  as  to  include 
whole  numbers  of  equivalent  weights,  will  not  all  occupy  exactly  equal  volumes. 
They  represent  exactly  equal  numbers  of  molecules,  and  the  slight  differences  in 
the  closeness  with  which  the  molecules  are  packed  (p.  234)  will  cause  the  values  of 
the  molar  volumes  to  differ  from  gas  to  gas.  The  values  of  the  volume  actually 
containing  equal  numbers  of  molecules  are  as  follows:  Hydrogen,  22.40;  oxygen, 
22.39;  nitrogen,  22.45;  chlorine,  22.01;  hydrogen  chloride,  22.23;  carbon 
dioxide,  22.26;  water,  22.39;  mercury,  22.55.  The  average  value  of  this  volume 
in  the  case  of  the  more  nearly  perfect  gases  is  22.4  liters,  and  this  is,  therefore, 
the  number  which  we  have  used  in  our  definition. 

*  A  common  question  is:  Do  not  molecules  of  different  substances  differ  in 
size,  and  will  not  the  numbers  required  to  fill  the  G.M.V.  therefore  be  different? 
The  answer  is  that  the  molecules  are  all  so  small  compared  with  the  spaces 
between  them  (at  760  mm.)  that  the  distances  from  surface  to  surface  are  practi- 
cally the  same  as  from  center  to  center.  A  G.M.V.  of  oxygen,  when  liquefied, 
gives  less  than  32  c.c.  of  liquid  oxygen,  or  less  than  1/700  of  the  volume  as  gas. 
It  is  only  when  gases  are  so  severely  compressed  that  the  nearness  of  the  molecules 
to  one  another  approaches  that  found  in  the  liquid  condition  that  the  effects  of  the 
bulk  of  the  molecules  become  conspicuous,  and  a  difference  in  the  behavior  of 
different  gases  becomes  noticeable.  But  in  the  kind  of  chemical  work  discussed  in 
this  chapter,  pressures  over  one  atmosphere  are  not  used. 


238  INORGANIC  CHEMISTRY 

The  Number  of  Molecules  in  a  Mole  or  in  22.4  Liters.  — 

The  molecular  weight  or  mole  of  a  substance  is  not  the  weight  of  a 
single  molecule.  It  is  only  the  relative  weight  of  the  molecule  of  the 
substance.  It  is,  however,  the  weight  in  grams  of  a  fixed  number 
of  molecules,  for  22.4  liters  (or  any  other  volume)  contains  equal 
numbers  of  molecules  of  different  gases.  The  actual  number  has 
been  determined.  Thus  Jean  Perrin  found  values  by  several  experi- 
mental methods  which  ranged  between  5.9  X  1023  (that  is,  59  fol- 
lowed by  22  ciphers)  and  6.9  X  1023.  Rutherford,  using  an  entirely 
different  plan,  obtained  5.7  X  1023  for  the  gas  helium.  This  value 
was  obtained  by  observing  that  0.46  cubic  millimeters  of  helium 
(0°  and  760  mm.)  were  given  off  by  1  gram  of  radium  (q.v.)  every 
24  hours.  Now  the  molecules  of  helium,  when  they  strike  a  screen 
covered  with  zinc  sulphide,  produce  flashes  of  light.  By  counting 
the  flashes  appearing  in  a  certain  restricted  area  in  one  minute,  and 
multiplying  by  the  proper  factor,  to  get  the  total  number  emitted 
on  all  sides  during  that  time,  and  further  multiplying  by  1440  min- 
utes (=24  hours),  the  number  of  molecules  filling  0.46  c.  mm.  was 
obtained.  From  the  result,  the  numbers  in  1  c.c.  and  in  22.4  1.  were 
calculated.  The  value  at  present  accepted  as  most  accurate  is 
that  obtained  by  R.  A.  Millikan  of  the  University  of  Chicago,  by 
the  use  of  a  still  different  method,  namely  6.06  X  1023  (or  60602i). 

Accepting  one  of  these  values,  e.g.,  that  of  Millikan,  we  can  evi- 
dently calculate  the  weight  of  a  single  molecule  of  any  gaseous  or 
volatile  substance.  Thus,  a  single  molecule  of  oxygen  must  weigh 
32/6.06  X  1023  g.,  or  5.3  X  1Q-23,  or  0.02253  g. 

The  Chemical  Molecule.  —  It  will  be  noted  that  we  get  defi- 
nite, quantitative  information  about  individual  molecules  by  the 
study  of  gases.  We  shall  find,  later,  that  we  can  also  obtain  similar 
information  about  molecules  of  substances  in  dilute  solutions,  be- 
cause, in  that  condition,  the  molecules  are  separated  from  one  an- 
other and  scattered  throughout  a  considerable  space,  much  as  they 
are  in  gases  (p.  184).  Chemical  molecules  are  the  units  of  which 
gases,  and  bodies  in  dilute  solution,  are  aggregates.  We  have 
means  of  comparing,  roughly,  at  least,  the  dimensions  of  the  physical 
units  in  gases  and  liquids,  and  find  that  in  the  liquid  state  the  mole- 
cules are  often  multiples  of  the  gaseous  ones  (c/.  water,  p.  202). 
The  gaseous  molecules  are,  therefore,  our  standard  in  chemistry. 

Many  authors  discuss  the  breaking  up  of  a  solid,  like  salt  or  glass,  until 
particles  are  obtained  which  cannot  be  further  split,  They  use  this  in  the 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          239 

attempt  to  make  clear  that  there  is  a  limit  to  pulverization,  reached  when  the 
material  has  been  reduced  to  single  molecules.  Now,  this  process  is  purely 
imaginary,  and  cannot  actually  be  carried  out.  Chemistry  is  an  experimental 
science,  and  only  the  real  experiments  which  lead  to  a  given  .result  should  be 
described.  This  statement  about  pulverization  is  disingenuous,  to  say  the  least, 
for  there  are  no  facts  behind  it.  Aside  from  that  objection,  no  conditions  are 
usually  stated  that  would  prevent  one  from  dividing  the  molecules  into  atoms, 
and  the  atoms  into  electrons  and  positively  charged  nuclei  (see  Chap.  XXII). 
Since  the  process  is  usually  stated  to  be  imaginary,  there  is  no  particle  so  small 
but  that  one  can  imagine  it  being  split  into  two,  or  even  twenty  parts,  and  there 
is  thus  no  way  of  stopping  the  imaginary  splitting  at  the  right  point.  An  active 
imagination  could  split  salt  into  small  fractions  of  electrons,  although  we  have  no 
experimental  knowledge  of  the  existence  of  such  fragments.  The  imagination  is 
of  great  value  in  suggesting  possibilities,  to  be  tested  experimentally,  but  it  cannot 
be  used,  alone,  for  ascertaining  facts,  and  it  is  crime  against  the  spirit  of  true 
science  to  suggest  the  contrary  to  a  naturally  confiding  reader. 


ATOMIC  WEIGHTS 

Determination  of  the  Atomic  Weight  of  Each  Element.  — 

If  the  paragraphs  dealing  with  combining  weights  are  now  re-read 
(pp.  61-63),  it  will  be  found  that  the  foundations  for  a  system  of 
weights  was  worked  out,  but  that  no  basis  for  definitely  fixing  the 
individual  values  was  discovered.  At  the  time,  the  only  informa- 
tion we  had  was  obtained  by  analyzing  compounds  and  reasoning 
about  the  results,  and  evidently  something  more  was  needed  for  the 
absolute  determination  of  the  values.  Thus  on  p.  66,  it  is  pointed 
out  that  the  equivalent  weight,  or  any  multiple  of  it  by  an  integer, 
will  serve  for  expressing  the  proportions  used  by  the  element  in  com- 
bining with  other  elements.  We  are  now  approaching  the  question 
of  units,  in  which  to  express  combining  proportions,  from  a  different 
viewpoint.  We  were  then  assigning  numbers  for  the  quantities  of 
the  constituent  elements  of  a  compound  (such  as  iron  :  oxygen  :  : 
111.68  :48,  p.  13)  without  any  consideration  of  the  magnitude  of 
the  total  weight  of  the  constituents.  At  that  time,  we  had  no 
reason  before  us  to  indicate  that  this  total  might  require  considera- 
tion. We  now  start  by  determining  and  assigning  the  total  weight 
of  the  compound,  and  it  is  our  next  task  to  consider  the  subdi- 
vision of  this  total  amongst  the  constituents.  Evidently,  if  the  unit 
quantity  of  the  compound  has  been  properly  chosen,  it  must  be 
sub-divisible  into  one  or  more  unit  quantities,  of  suitable  dimen- 
sions, of  each  element  in  the  compound. 


240  INORGANIC  CHEMISTRY 

To  determine  the  atomic  weights,  the  plan  of  procedure  is  per- 
fectly simple.  In  the  preceding  section  we  settled  upon  the  chemical 
unit  quantity  of  each  substance.  This  is  the  quantity  which,  in  the 
gaseous  condition,  would  fill  the  G.M.V.  (22.4  liters)  at  0°  and  760 
mm.  Now,  we  seek  the  chemical  unit  quantities  of  the  elements 
combined  in  each  substance.  Evidently  the  logical  and  consistent 
plan  must  be  to  take  the  amount  of  each  substance  which  fills  the 
G.M.V.  and  find  out  how  much  of  each  element  present  is  contained 
in  this  unit  amount  of  the  substance.  In  other  words>  to  put  the 
matter  concretely,  we  imagine  ourselves  filling  the  cube  (Fig.  93) 
with  one  compound  after  another,  and  in  each  case  determining  by 
analysis  the  weight  of  each  constituent  element  present  in  a  cube- 
full  of  the  substance.  To  carry  out  this  plan,  two  experimental 
operations  are  necessary  with  each  substance: 

First  we  determine  the  density,  and  this  gives  us  the  gram-molec- 
ular weight,  i.e.,  the  amount  filling  the  cube.  This  shows  the  rela- 
tive weight  of  a  molecule  of  the  substance,  as  compared  with  that  of 
one  molecule  of  oxygen. 

Second  we  analyze  the  substance,  and  this  gives  us  the  quantity 
of  each  constituent  in  the  total  gram-molecular  weight,  i.e.,  in  the 
material  filling  the  cube.  This,  in  turn,  shows  the  weight  of  the 
quantity  of  each  element  present  in  each  molecule,  relative  to 
the  weight  of  a  molecule  of  oxygen. 

For  example,  the  cube  holds  36.468  g.  of  hydrogen  chloride,  and 
this  amount,  when  decomposed,  yields  1.008  g.*  of  hydrogen  and 
35.46  g.  of  chlorine. 

Finally,  to  determine  the  best  unit  weight  for  a  given  element, 
we  repeat  the  two  foregoing  operations  with  as  many  different  com- 
pounds of  the  element  as  possible,  and  then  we  examine  the  various 
quantities  of  the  element  found  in  the  G.M.V.  of  the  various  com- 
pounds. From  inspection  of  these  quantities  we  quickly  select  the 
value  of  which  all  are  multiples,  by  unity  or  some  integral  number. 
This  value  for  the  unit  weight  is  the  one  accepted  as  the  atomic 
weight.  This  is  the  weight  of  one  atom  of  the  element,  compared 
with  the  weight  of  a  molecule  of  oxygen,  and  molecules  containing 

*  It  will  be  observed  that  if  the  unit  for  molecular  weights  had  been  less  than 
the  number  of  molecules  in  32  g.  of  oxygen,  then  an  equal  number  of  molecules  of 
hydrogen  chloride  would  have  contained  less  than  1.008  g.  of  hydrogen,  and  the 
atomic  weight  of  this  element  would  then  have  been  less  than  unity.  Theoreti- 
cally, any  value  for  hydrogen  would  serve  the  purpose,  but  there  are  arithmetical 
conveniences  in  having  one  value  in  the  list  close  to  unity. 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS 


241 


more  than  this  proportion  contain  two,  three,  or  more  atoms  of  the 
element. 

For  example,  if  we  are  seeking  the  atomic  weight  of  chlorine,  we 
set  down  the  result  for  hydrogen  chloride  just  given.  Then  we  take 
another  compound  of  chlorine,  say  phosphorus  oxychloride  (a  liquid) . 
We  determine  the  weight  of  a  measured  volume  of  its  vapor,  at  a 
properly  chosen  temperature  and  pressure,  and  the  result  gives  us,  by 
calculation,  the  molecular  weight,  viz.  153.38.  That  is,  153.38  g.  of 
the  substance  would  fill  the  cube,  if  it  could  be  kept  as  vapor  at  0° 
and  760  mm.  And  this  amount  of  the  substance  contains  31  g.  of 
the  element  phosphorus,  16  g.  of  the  element  oxygen,  and  106.38  g. 
of  the  element  chlorine.  We  then  continue  the  processes  described, 
using  all  the  volatile  compounds  of  chlorine.  The  involatile  com- 
pounds (like  common  salt)  must  be  set  aside,  for  they  cannot  be 
vaporized,  and  therefore  their  molecular  weights  cannot  be  deter- 
mined. When  we  have  studied  as  many  compounds  as  possible  in 
this  way,  we  find  that  there  are  different  quantities  of  chlorine  in 
our  list,  but  they  are  all  integral  multiples  of  35.46  g.  In  phosphorus 


Substance. 

Molar 
Weight. 

Weights  of  Constituents  in 
Molar  Weight. 

Hydrogen. 

Chlorine. 

| 

1 

32 

16 
160 

'ie 

Phosphorus. 

Carbon. 

Mercury. 

Molecular 
Formula. 

Hydrogen  chloride     
Chlorine  dioxide 
Phosphorus  trichloride     .... 
Phosphorus  oxychloride    .... 
Phosphoric  anhydride           .    .    . 

36.46 
67.46 
137.38 
153.38 
284 
34 
18 
'16 
26 
44 
30 
60 
235.46 
270.92 

1 

35.46 

35.46 
106.38 
106.38 

HC1 
C102 
PC13 
POCla 
P4010 
PH3 
H2O 
CH4 
C2H2 
C3H8 
CH2O 
C2H402 
HgCl 
HgCl2 

31 
31 
124 
31 

Phosphine                     .           ... 

3 
2 

4 
2 
8 
2 
4 

12 
24 
36 
12 
24 

200 
200 

Water                        

Miethane                       

Acetylene             

Propane             

Formaldehyde     

35^46 
70.92 

16 
32 

Acetic  acid    

Mercurous  chloride    .    .    .  '.    .    . 

Mercuric  chloride  . 

oxychloride,  for  example,  the  quantity  was  106.38,  or  3  X  35.46. 
Hence  35.46  g.  can  be  taken  as  the  unit  quantity,  the  atomic  weight 
of  the  element  chlorine.  This  is  the  relative  weight  of  an  atom  of 


242  INORGANIC  CHEMISTRY 

chlorine,  as  compared  with  the  weight  of  a  molecule  of  oxygen,  when 
the  value  32  is  assigned  to  the  latter.* 

In  the  preceding  table  a  few  sample  results  of  the  process  just  out- 
lined are  given.  The  first  column  contains  the  molar  weight,  i.e.,  the 
weight  of  the  substance  which  occupies  the  G.M.V.  cube.  In  the 
other  columns  are  entered  the  weights  of  the  various  elements  which 
together  make  up  the  total  molar  weight.  To  simplify  the  numbers, 
the  value  of  1  is  used  for  hydrogen,  instead  of  1.008. 

To  contain  similar  data  for  all  the  volatile  compounds  of  every 
known  element,  a  huge  table,  of  which  this  might  be  a  small  corner, 
would  be  required.  With  such  a  table  at  hand  the  atomic  weight 
of  each  element  could  promptly  be  picked  out.  Thus,  in  the  carbon 
column  it  would  be  found  that  all  the  weights  of  carbon  were  either 
12  or  integral  multiples  of  12,  and  this  is  therefore  the  atomic  weight 
of  carbon.  Similarly  the  atomic  weight  of  oxygen  is  16, f  of  phos- 
phorus 31,  of  mercury  200  (see  Exercise  4). 

When  the  atomic  weights  have  finally  been  selected,  we  can  go 
through  the  table  and  change  all  the  numbers  into  multiples  of  the 
chosen  atomic  weights.  Thus,  for  70.92  we  write  2  X  35.46,  and 
for  106.38  we  write  3  X  35.46,  and  so  forth.  The  reader  should  pre- 
pare such  a  modification  of  the  table.  With  this  new  form  of  the 
table  before  us,  we  can,  finally,  replace  the  atomic  weights  by  the 
symbols  which  stand  for  them,  writing,  for  35.46,  Cl,  for  2  X  35.46, 
Cl2,  and  so  forth.  The  results  of  doing  this  in  each  line,  i.e.,  for 
each  substance,  are  collected  at  the  ends  of  the  lines  in  the  last  col- 
umn of  the  table.  The  reader  should  himself  repeat  the  substitu- 
tions of  the  symbols,  and  so  verify  the  formulae  given.  These 
formulas,  since  they  are  based  on  the  molecular  weights,  in  such  a 
way  that  when  the  numerical  values  are  substituted  for  the  symbols 
the  total  restores  to  us  the  molecular  weight,  are  called  molecular 
formulae. 

The  mode  of  derivation  makes  all  the  quantities  of  any  one  ele- 
ment in  the  above  table  integral  multiples  of  the  smallest.  Of 
course,  an  element  might  be  found  of  which  no  volatile  compound 

*  It  should  be  noted  that  there  is  another  unit  quantity  of  chlorine,  namely  the 
molecular  weight,  or  weight  of  the  G.M.V.  of  the  substance.  This  is  the  unit 
quantity  of  free  chlorine.  But  we  are  dealing  now  with  compounds,  and  propor- 
tions in  combination,  so  that  free,  uncombined  chlorine,  and  other  elements  in 
free  condition  do  not  interest  us  at  present,  and  will  be  taken  up  later. 

t  The  difference  between  the  unit  quantity  of  oxygen  in  compounds  (namely, 
16)  and  the  unity  quantity  of  free  oxygen  (32)  will  be  discussed  presently, 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          243 

containing  the  unit  weight  was  yet  known.  If  methane  and  formal- 
dehyde were  still  undiscovered,  and  only  the  other  three  compounds 
containing  carbon  were  available,  we  should  then  take  the  greatest 
common  measure  of  24  and  36,  namely  12,  as  the  unit. 

The  chemical  unit  weight  or  atomic  weight  of  an  element  may 
thus  be  defined  as:  The  G.C.M.  (greatest  common  measure)  of  the 
weights  of  that  element  found  in  the  molar  weights  of  all  its  volatile 
compounds  so  far  as  these  have  been  examined. 

Atomic  Weights  and  Equivalents.  —  It  will  now  be  seen  why 
the  equivalents  (p.  63)  were  multiplied  by  various  integers  in 
making  the  chemical  units.  The  equivalent  of  carbon  was  3.  That 
is  to  say,  carbon  and  oxygen  combine  in  the  ratio  3  :  8  (in  carbon 
dioxide),  and  carbon  and  hydrogen  in  the  ratio  3  :  1.008  (in  methane). 
But  there  is  no  compound  of  carbon  whose  molecular  weight  con- 
tains less  than  12  parts  of  the  element.  It  would  thus  lead  to  need- 
less complication  to  take  3  as  the  unit  amount  of  carbon,  for  every 
molecule  would  then  contain  four  units,  or  some  multiple  of  four, 
and  every  formula  C4  or  some  multiple  of  C4.  We  choose  the  largest 
units  of  combining  weight  that  we  can,  in  order  that  the  coefficients 
may  be  the  smallest  possible,  and  the  resulting  formulae  the  simplest 
possible.  Naturally  the  actual  ratios  remain  the  same.  Thus,  for 
carbon  dioxide  the  ratio  3  :  8  is  replaced  by  12  :  32,  or  12  :  2  X  16, 
or  C  :  20,  which  has  the  same  value. 

Since  the  atomie  weight  is  thus  always  a  multiple  of  the  equivalent 
weight  (by  unity,  or  some  other  integer) ,  it  might  also  be  defined  as : 
The  largest  integral  multiple  of  the  equivalent  which  can  be  contained 
in  the  molecular  weights  of  all  the  volatile  compounds  of  the  element. 
The  complete  list  of  accepted  atomic  weights  is  printed  on  the  inside 
of  the  cover  at  the  back  of  this  book. 

The  conclusion  that  C  =  12  is  the  most  convenient  value  is 
reached  also  by  studying  the  chemical  interactions  of  carbon  com- 
pounds. If  the  symbol  C  stood  for  three  parts  (the  equivalent)  of 
carbon,  methane  would  receive  the  formula  CH,  the  proportion  of 
its  constituents  by  weight  being  3  of  carbon  and  1  of  hydrogen. 
But,  when  we  mix  this  gas  with  chlorine  and  expose  the  mixture  to 
sunlight,  no  less  than  four  different  compounds  are  produced  (cf.  p. 
225).  With  C  =  3,  their  formula  would  be  C4H3C1,  C2HC1,  C4HC13, 
and  CC1,  the  carbon  being  univalent.  The  relations  of  these  sub- 
stances are  much  simplified  when  we  change  to  molecular  formulae 
and  substitute  C  =  12  for  C*  =  12,  making  the  carbon  quadrivalent. 


244  INORGANIC  CHEMISTRY 

We  then  perceive  that  we  are  displacing  successively  the  four  hydro- 
gen units  in  one  molecule,  and  that  the  substances  are  CH4,  CH3C1, 
CH2C12,  CHC13,  and  CCU.  The  whole  prodigious  growth  of  the 
chemistry  of  the  compounds  of  carbon,  which  has  taken  place  during 
the  last  half  century,  has  been  the  result  of  the  employment  of  this 
seemingly  slight  improvement  by  Gerhardt  and  Laurent,  and  by 
Williamson,  and  its  enforcement  and  extension  by  Kekule  and  Cou- 
per,  independently  of  one  another,  in  1858.  Thus,  quite  aside  from 
the  molecular  theory  and  Avogadro's  addition  to  it,  we  have  now 
found  ample  independent  justification  for  the  multiplication  of  the 
equivalents  by  integers  and  for  the  conception  of  valence  which 
results  from  this. 

Further  Comments  on  Atomic  Weights.  —  The  fact  that 
all  the  numbers  in  any  one  column  of  the  table  (p.  241),  turn  out  to 
be  even  multiples  of  a  single  number  need  not  seem  mysterious. 
The  molecule  of  every  compound  containing  chlorine  must  contain 
one,  two,  three,  or  some  other  whole  number  of  chlorine  atoms,  for 
chlorine  atoms,  like  other  atoms,  do  not  furnish  fractions  of  atoms 
in  any  cases  of  combination.  Now,  the  weight  of  chlorine  in  60602i 
atoms,  assuming  one  atom  of  chlorine  to  each  molecule  in  22.4  liters, 
must  be  35.46  g.  Hence,  if  the  weight  of  chlorine  in  22.4  liters 
(60602i  molecules)  of  the  compound  differs  from  35.46  g.,  it  can  do 
so  only  because  there  are  two  atoms  of  chlorine  per  molecule,  giving 
2  X  35.46  g.,  or  three  atoms  giving  3  X  35.46  g.  of  chlorine,  and  so 
forth.  Thus  the  quantities  of  chlorine  in  the  G.M.V.  of  all  com- 
pounds of  chlorine  must  be  multiples  of  35.46  by  unity  or  some 
other  integer. 

It  is  hardly  necessary  to  add  that  the  atomic  weights,  found  as 
described  above,  are  equally  serviceable  in  expressing  the  composi- 
tions of  compounds  which  are  not  volatile.  The  atoms  in  non- 
volatile compounds  are  identical  in  properties  with  the  atoms  of  the 
same  elements  in  volatile  compounds.  If  an  element  gives  no  vola- 
tile compounds,  other  methods  of  fixing  its  atomic  weight  are  avail- 
able (see  Dulong  and  Petit's  law,  p.  245). 

Although  in  this  section,  as  well  as  elsewhere,  we  have  empha- 
sized the  fact  that  atoms  are  not  divided  into  parts,  this  must  not 
be  taken  to  mean  that  atoms  are  incapable  of  being  broken  up.  It 
means  only  that  in  ordinary  chemical  changes,  the  atoms  combine 
and  separate  as  wholes.  Indeed,  we  now  know  that  the  atom  of 
radium  (q.v.)  gives  off  atoms  of  helium,  and  leaves  an  atom  of  lead, 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          245 

and  that  the  atoms  of  one  or  two  other  elements  disintegrate  in  a 
similar  way.  Some  day  means  of  breaking  up  any  or  all  kinds  of 
atoms  may  be  discovered. 

Advantages  of  Atomic  Weights.  —  Although  the  method  of 
selecting  atomic  weights  involves  rather  complex  reasoning,  these 
weights  repay  the  trouble,  because  they  represent  the  relative 
weights  of  the  atoms  themselves.  They  are  thus  much  more  val- 
uable in  helping  us  to  understand  chemical  behavior  and  in  enabling 
us  to  classify  the  phenomena  of  chemistry  than  would  be  any  other 
units  of  weight  we  might  have  chosen.  The  following  are  some  of 
the  advantages  they  offer: 

1.  The  atomic  weight  of  an  element  has  but  one  value,  and  this 
value  is  definitely  determinable. 

2.  The  atomic  weight  of  an  element  has  a  valence  (p.  130),  while 
equivalents  are  equi-valent.     While  valence  is  a  helpful  conception 
in  all  branches  of  chemistry.,  organic  chemistry  is  especially  indebted 
to  the  conception  of  the  quadrivalence  of  carbon  for  much  of  its  de- 
velopment and  most  of  its  organization.     The  full  illustration  of 
this  point  is  beyond  the  limits  of  the  present  book. 

3.  The  periodic  system  (q.v.),  the  basis  of  a  plan  for  classifying 
the  properties  of  all  chemical  substances,  is  founded  upon  the  atomic 
weights. 

4.  Dulong  and  Petit's  law  is  based  upon  atomic  weights.     This 
law  furnishes  also  an  alternative  means  of  determining  atomic  weights 
that  has  frequently  rendered  valuable  service,  and  on  this  account 
forms  the  subject  of  the  next  section. 

Dulong  and  Petit's  Law,  an  Alternative  Means  of  Deter- 
mining Atomic  Weights.  —  It  was  first  pointed  out  (1818)  by 
Dulong  and  Petit,  of  the  Ecole  Polytechnique  in  Paris,  that  when 
the  atomic  weights  of  the  elements  were  multiplied  by  the  specific 
heats  of  the  simple  substances  in  the  solid  condition,  the  products 
were  approximately  the  same  in  all  cases.  In  other  words,  the  spe- 
cific heats  are  inversely  proportional  to  the  magnitudes  of  the  atomic 
weights.  The  table,  in  which  round  numbers  have  been  used  for 
the  atomic  weights,  shows  that  the  product  lies  usually  between  6 
and  7,  averaging  about  6.4; 


246 


INORGANIC   CHEMISTRY 


Element. 

Atomic 
Wt. 

Sp.  Ht. 

Prod- 
uct. 

Element. 

Atomic 
Wt. 

Sp.  Ht. 

Prod- 
uct. 

Lithium     .    . 

7 

.94 

6.6 

Iron     .... 

56 

112 

6  3 

Sodium      .    . 

23 

.29 

6.7 

Zinc    . 

65  4 

093 

6  1 

Magnesium  . 

24.3 

.245 

6.0 

Bromine  (Solid) 

80 

.084 

6.7 

Silicon    .    .    . 

28.3 

.16 

4.5 

Gold    . 

197 

032 

6  3 

Phosphorus 
(Yellow) 

31 

.19 

5.9 

Mercury 
(Solid) 

200 

.0335 

6.7 

Calcium     .    . 

40 

.170 

6.8 

jUranium    . 

238 

.0276 

6.6 

Another  way  of  expressing  this  law  will  give  it  greater  chemical 
significance.  The  specific  heats  are  the  amounts  of  heat  required  to 
raise  one  gram,  that  is  one  physical  unit,  of  each  element  through 
one  degree.  When  we  multiply  this  by  the  atomic  weight,  we 
obtain  the  amount  of  heat  required  to  raise  one  gram-atomic  weight 
of  the  element,  that  is,  one  chemical  unit,  through  one  degree.  The 
values  of  this  product  are  approximately  equal.  Since  there  are 
equal  numbers  of  atoms  in  one  gram-atomic  weight  of  each  element, 
it  follows  that :  Equal  amounts  of  heat  raise  equal  numbers  of  atoms  of 
all  elements  in  the  solid  form  through  equal  intervals  of  temperature. 

The  conspicuous  exceptions  occur  among  the  elements  with  low  atomic 
weights  only.  The  products  for  four  of  these,  the  atomic  weights  being  given  in 
parentheses,  are  as  follows:  Glucinum  (9),  3.7;  boron  (11),  2.8;  carbon  (12),  1.7; 
and  silicon  (28.4),  4.5.  Further  investigation  shows,  however,  that  in  the  case  of 
precisely  these  elements  the  value  found  for  the  specific  heat  varies  very  markedly 
with  the  temperature  at  which  the  specific  heat  is  measured.  Their  specific  heats 
become  rapidly  greater  the  higher  the  temperature.  Thus  at  985°  the  specific  heat 
of  the  diamond  (carbon)  is  0.45,  which  when  multiplied  by  the  atomic  weight 
gives  the  product  5.5.  So  that  even  the  exceptional  elements  tend  to  come  into 
line  when  the  specific  heat  is  measured  at  higher  temperatures  —  and  Dulong 
and  Petit's  statement  does  not  limit  our  choice  to  any  one  temperature. 

It  will  be  seen  at  once  that  although  the  law  of  Dulong  and  Petit 
is  purely  empirical,  it  may  nevertheless  be  used  for  fixing  the  atomic 
weight  of  an  element  of  which  no  volatile  compounds  are  known.  We 
can  always  measure  the  equivalent  with  considerable  exactness,  and, 
when  this  has  been  multiplied  by  the  specific  heat  of  the  free  sub- 
stance, we  can  see  at  a  glance  what  integral  factor  will  raise  the 
product  to  the  neighborhood  of  6.4.  For  example,  analysis  shows  us 
that  in  calcium  chloride  the  proportion  of  chlorine  to  calcium,  using 
the  known  atomic  weight  of  chlorine  as  one  term  of  the  proportion, 
is  35.46  :  20.  If  calcium  is  univalent,  20  is  its  atomic  weight.  If  it 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          247 

is  bivalent,  two  units  of  chlorine  are  combined  with  40  parts  of  cal- 
cium, and  40  is  its  atomic  weight.  If  it  is  trivalent,  three  units  of 
chlorine  are  united  with  60  parts  of  calcium,  etc.  All  we  learn  in 
reference  to  the  atomic  weight  of  calcium  from  this  analysis  is  that 
its  value  is  20  or  some  integral  multiple  of  20.  Nor  can  we  fix  the 
upper  limit,  for  we  are  unable  to  obtain  the  weight  of  a  known 
volume  of  calcium  chloride  vapor  and  so  determine  the  molecular 
weight.  But  the  specific  heat  of  solid  calcium  being  0.170,  we  mul- 
tiply this  number  by  20,  and  get  the  product  3.4.  This  is  only  half 
large  enough,  so  we  find  that  40  is  the  value  for  the  atomic  weight 
of  calcium.  The  product  is  then  6.8,  which  agrees  fairly  well  with 
the  average  for  other  elements.  We  decide,  therefore,  that  the 
symbol  Ca  shall  represent  forty  parts  by  weight.  The  formula  of 
calcium  chloride  is  therefore  CaCl2,  and  calcium  is  bivalent. 

It  will  be  seen  that  this  does  not  supply  us  with  a  method  of  ascertaining 
chemical  unit  weights  independently  of  any  chemical  experiment.  We  cannot 
measure  the  specific  heat  and  use  the  quotient  from  division  of  this  number  into 
6.4,  for  we  do  not  know  in  advance  that  the  product  for  the  element  will  have 
exactly  this  value.  It  may  be  below  6,  or  it  may  be  as  high  as  7.  In  the  case  of 
calcium,  for  example,  6.4  -r-  0.17  =  37.65.  Now  37.65  is  5  per  cent  below  the  real 
value  of  the  chemical  unit,  and  even  the  roughest  measurement  of  a  chemical 
combining  weight  need  never  be  more  than  1  per  cent  in  error.  Hence  the  atomic 
weight  must  be  founded  upon  the  determination  of  the  equivalent,  which  can 
be  measured  with  accuracy.  The  rule  discussed  in  this  section  can  be  used  only 
to  ascertain  what  multiple  of  the  equivalent  shall  be  accepted  as  the  atomic  weight 
after  the  equivalent  itself  has  been  measured  with  care.  In  other  words,  this  is  a 
method  of  adjusting  the  result  of  chemical  experimentation,  and  cannot  supersede 
it  altogether. 

The  existence  of  the  law  of  Dulong  and  Petit  and  the  periodic  law,  together 
with  the  services  of  structural  formulae  to  organic  chemistry,  all  demonstrate  that 
atomic  weights  are  of  vastly  greater  significance  in  the  science  than  are  equivalent 
weights.  And  there  are  other  immense  ranges  of  facts,  aside  from  those  covered 
by  these  conceptions,  which  are  all  dependent  upon  the  atomic  weights.  That 
almost  the  whole  systematization  that  has  been  secured  in  chemistry  should  thus 
center  in  this  one  point,  furnishes  the  strongest  circumstantial  evidence  that 
Avogadro's  law  is  hi  accord  with  the  facts.  This  independent  inductive  evidence 
in  favor  of  Avogadro's  principle  is  especially  worth  noting  because  the  deduction 
of  the  principle  from  the  data  of  the  kinetic-molecular  theory  is  not  absolutely 
rigid.  It  involves  certain  assumptions  which,  while  they  are  plausible  enough, 
are  still  assumptions. 

Historical.  —  The  idea  that  matter  is  composed  of  small  par- 
ticles is  a  very  ancient  one.  Not  even  Dalton  (1805),  however, 


248  INORGANIC  CHEMISTRY 

although  he  used  this  conception  continually  as  a  means  of  thought 
about  chemical  and  physical  phenomena,  made  any  distinction  be- 
tween atoms  and  molecules.  This  distinction  was  introduced  later 
by  Avogadro  (1811).  Yet,  without  this  refinement,  continual 
thought  of  the  behavior  of  matter  as  consisting  in  transactions  be- 
tween small  particles  led  Dalton  to  see  that  it  was  probable  that 
individual  unit  weights  for  each  element  must  exist.  The  discovery 
that  they  did  exist  soon  followed.  The  numbers  which  Dalton 
actually  gave  out,  aside  from  the  considerable  experimental  inaccu- 
racies attached  to  them,  were  often  equivalents  and  not  modern 
atomic  weights.  It  was  after  the  publication  of  Dalton's  ideas  that 
Gay-Lussac  discovered  the  law  of  combining  volumes  (1808),  and 
until  this  law  was  discovered  there  was  no  criterion  for  fixing  the 
values  of  the  atomic  weights.  Gay-Lussac,  at  the  end  of  his  paper, 
pointed  out  that  his  discovery  formed  an  important  confirmation 
of  Dalton's  views.  Strange  to  say,  Dalton  himself  refused  to  accept 
Gay-Lussac's  law,  and  so  rejected  the  very  means  by  which  his  own 
principle  of  chemical  unit  weights  came  eventually  to  be  acknowl- 
edged as  one  of  the  foundation  stones  of  the  science.  On  the  other 
hand,  Dalton's  fellow  countrymen  and  contemporaries  accepted  the 
principle  of  unit  weights,  but  rejected  the  atomic  hypothesis  by  the 
help  of  which  Dalton  had  reached  them !  Thus  Sir  Humphry  Davy 
called  them  " proportions"  instead  of  atomic  weights,  and  Wollaston 
preferred  the  word  "  equivalents." 

Gerhardt  and  Laurent,  and  Williamson  used  Avogadro's  law 
long  before  1858,  and  employed  correct  atomic  weights,  so  far  as 
the  elements  possessed  volatile  compounds.  But  there  was  much 
difference  of  opinion  about  the  best  atomic  weights  for  the  other 
elements,  and  some  difference  even  in  regard  to  elements  like  oxygen. 
Thus,  many  still  wrote  the  formula  of  water  HO,  where  O  =  8.  It 
was  only  after  Cannizzaro  (1858)  demonstrated  the  value  of  Dulong 
and  Petit's  law  (discovered  forty  years  before)  for  settling  the 
atomic  weights  of  elements,  like  calcium  and  iron,  which  gave  no 
easily  volatile  compounds,  that  rapid  progress  towards  complete 
agreement  was  made. 

Many  chemists  have  contributed  to  the  determination  and  revi- 
sion of  the  atomic  weights.  The  Swedish  chemist,  Berzelius,  devoted 
many  years  to  the  accurate  measurement  of  combining  proportions. 
Stas,  a  Belgian  (1860-1870),  made  a  number  of  determinations 
with  great  exactness.  Morley's  (1895)  value  for  combining  propor- 
tions of  hydrogen  and  oxygen  alone  represented  twelve  years  of 


MOLECULAR   WEIGHTS  AND  ATOMIC  WEIGHTS          249 

work.     T.  W.  Richards  of  Harvard  University  has  carried  many  of 
the  values  to  a  higher  degree  of  accuracy. 

MOLECULAR  FORMULAE  • 

Molecular  Formulae  of  Compounds.  —  If  the  molar  formulae 
in  the  table  (p.  241)  be  examined  it  will  be  observed  that  several  are 
not  in  their  simplest  terms.  Thus,  the  formula  of  acetylene  is  C2H2. 
The  formula  CH  would  represent  the  composition  of  the  substance 
equally  well,  for  12  :  1  is  the  same  as  24  :  2.  But  the  formula  CH 
gives  a  total  of  only  13,  while  C2H2  shows  the  total  weight  of  the 
molecule  to  be  26  and  records  for  us  therefore  the  weight  of  the  G.M.V., 
as  well  as  the  composition  of  the  substance.  We  shall  find  this 
additional  property,  peculiar  to  the  molecular  formula,  to  be  a  feature 
of  the  greatest  practical  value.  Some  of  the  practical  uses  of  this 
improvement  in  our  formulae  will  be  illustrated  in  this  chapter,  and 
there  is  an  example  of  one  of  them  in  the  table  itself.  Thus,  the 
molecular  formula  of  acetic  acid  is  C2H402,  and  not  the  simpler, 
identical  proportion  CH2O.  The  latter  is  the  molecular  formula  of 
a  totally  different  substance,  formaldehyde,  now  much  used  as  a  dis- 
infectant. The  vapor  of  this  substance  has  only  half  the  density  of 
acetic  acid  vapor,  and  this  fact,  recorded  in  the  formula,  helps  to  re- 
mind us  that  the  substances  are  different.  Still  another  substance  of 
the  same  composition  is  grape  sugar  (dextrose),  CeH^Oe.  In  addi- 
tion to  this  and  other  practical  advantages,  molecular  formulae  sat- 
isfy also  the  claim  of  logical  consistence.  If  the  symbols  represent 
the  atomic  weights,  the  formulae  should  be  constructed  so  as  to  rep- 
resent the  molecular  weights. 

Molecular  formulae  like  C2H2  and  C2H4O2  are  easily  interpreted 
in  terms  of  the  atomic  hypothesis.  C  represents  one  atom  of  carbon, 
and  H  one  atom  of  hydrogen.  But  there  is  no  reason  why  a  mole- 
cule of  acetylene  should  not  contain  two  atoms  of  each  kind.  Simi- 
larly, the  molecule  of  formaldehyde  contains  four  atoms  (CH2O), 
and  one  of  acetic  acid  eight  atoms  (C2H402),  and  one  of  dextrose 
twenty-four  atoms  (CeH^Oe),  although  the  relative  numbers  of  each 
kind  are  the  same.  Indeed,  this  hypothesis  helps  to  clear  the  matter 
up,  for  chemists  go  so  far  as  to  account  for  the  chemical  behavior  of 
the  substances  by  an  imagined  geometrical  arrangement  of  the 
atoms  in  their  molecules,  and  these  three  kinds  of  molecules  are 
supposed  to  differ  in  structure  as  well  as  in  the  number  of  atoms  they 
contain. 


250 


INORGANIC   CHEMISTRY 


The  Molecular  Weights  and  Formulse  of  Elementary  Sub- 
stances. —  The  following  table  gives  the  densities  of  some  ele- 
mentary substances,  including  those  of  which  the  substances  last 
discussed  are  compgunds.  The  first  column  shows  the  atomic 
weight,  which  in  each  case  is  the  minimum  weight  of  the  element 
found  in  a  G.M.V.  of  any  compound.  For  example,  16  g.  of  oxygen 


Atomic 
Weight. 

Sym- 
bol. 

Density, 
O  =  32. 

Density   Fac- 
torized. 

Formula 
of  Free 
Element. 

Oxvuen 

16  00 

o 

32.00 

2  X  16.00 

O2 

Hydrogen    

1.008 

H 

2.016 

2X  1.008 

H2 

Chlorine 

35  46 

Cl 

70.92 

2  X  35.46 

C12 

Phosphorus     
Mercury 

31.0 
200  6 

P 

Hg 

124.0 
200.6 

4X31.0 
1  X  200  .  6 

P4 

Hg 

Ozone   
Cadmium    
Potassium  
Sodium 

16.00 
112.4 
39.10 
23.00 

0 
Cd 
K 

Na 

48.00 
112.4 
39.10 
23.00 

3X  16.00 
IX  112.4 
1X39.10 
1  X23.00 

03 
Cd 
K 

Na 

Zinc 

65.37 

Zn 

65.37 

1  X65.37 

Zn 

and  35.46  g.  of  chlorine  are  the  weights  in  the  amounts  of  water 
vapor  and  hydrogen  chloride,  respectively,  which  fill  the  cube  (22.4 
liters).  The  symbol,  in  the  next  column,  stands  for  this  quantity 
and  occurs  in  many  formulae,  such  as  H2O  and  HC1.  It  represents 
the  combining  unit  or  atom.  In  the  third  column  is  given  the  den- 
sity of  the  free,  elementary  substance.  This  number  of  grams  of  the 
simple  substance  fills  the  G.M.V.  and  this  number  is  the  molecular 
weight.  It  shows  the  weight  of  the  molecule  relative  to  the  weights 
of  the  other  molecules  in  the  same  column,  and  to  the  weights  of 
the  atoms  in  the  first  column.  In  the  last  two  columns  are  given 
the  densities  resolved  into  multiples  of  the  atomic  weights  and  the 
corresponding  formulas. 

The  reader  cannot  fail  to  note  a  striking  peculiarity.  In  the  case 
of  chlorine  the  molecular  weight  is  70.92,  while  the  atomic  weight  is 
35.46.  With  hydrogen  and  oxygen,  also,  the  molecular  weight  con- 
tains two  atomic  weights.  Yet  this  is  not  a  general  rule,  for  with 
mercury  and  several  other  elements  the  molecular  and  atomic  weights 
are  alike,  while  with  phosphorus  the  molecular  is  four  times  the 
atomic  weight.  Evidently  there  is  no  rule,  and  each  element  has  to 
be  subjected  to  separate  experimental  study.  The  result  is  that  for 
free,  elementary  chlorine  we  use  the  molecular  formula  C12,  for  free 
hydrogen  H2,  for  elementary,  uncombined  oxygen  the  formula  O2.  For 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          251 

a  substance  like  phosphorus,  which  is  not  a  gas  and  is  not  often 
measured  as  a  vapor,  the  formula  P  is  commonly  employed  by 
chemists,  to  avoid  the  larger  coefficients  which  P4  introduces  into 
equations,  although  theoretically  the  latter  formula  would  be  the 
strictly  correct  one. 

The  case  of  oxygen  demonstrates  clearly  the  necessity  of  using 
molecular  formulae,  even  for  simple  substances.  The  table  shows 
two  substances  containing  nothing  but  oxygen.  Ozone  (q.v.)  has  a 
molecular  weight  48,  being  a  gas  exactly  one-half  heavier  than  ordi- 
nary oxygen.  Its  formula,  therefore,  is  O3,  while  that  of  oxygen  is  O2. 
Oxygen  and  ozone  are  entirely  different  chemical  individuals.  The 
latter  has,  for  example,  a  strong  odor  and  is  much  more  active. 
Thus  polished  silver  remains  bright  indefinitely  in  pure  oxygen,  but 
oxidizes  quickly  when  placed  in  ozone. 

To  avoid  a  common  error,  the  reader  should  note  that  to  learn  the 
atomic  weight  of  an  element,  we  do  not  measure  the  molecular  weight 
of  the  simple  substance.  The  molecular  weight  of  the  elementary 
substance  may  be  a  multiple  of  the  atomic  weight,  and  we  find  out 
whether  it  is  such  a  multiple  only  after  the  atomic  weight  has  been 
determined.  The  atomic  weight  is  the  unit  weight  used  in  com- 
pounds, and  can  be  ascertained  only  by  a  study  of  compounds.  •  The 
molecular  weight  of  the  free  element  gives  us  only  a  value  which  we 
know  must  be  a  multiple  of  the  atomic  weight,  by  1  or  some  other 
integer.  Mol.  Wt.  =  At.  Wt.  X  x,  where  x  is  1  or  some  other 
integer. 

Further  Discussion  of  the  Molecular  Formulae  of  Elemen- 
tary Substances.  —  Some  further  explanation  may  be  required,  to 
the  end  that  the  reader  may  be  reconciled  to  accepting  the  formulae 
Cl2,  O2,  and  so  forth.  In  the  first  place,  he  should  note  how  these 
formulae  arose.  If  we  accept  Avogadro's  law,  and  the  inference 
from  it  to  the  effect  that  the  weights  of  equal  volumes  of  gases  are 
in  the  same  ratio  as  the  weights  of  their  individual  molecules,  then 
we  cannot  escape  the  conclusion  to  which  measuring  the  relative 
densities  of  free  chlorine  and  hydrogen  chloride,  for  example,  leads. 
The  ratio  of  their  densities  is  70.92  :  36.46.  That  is  to  say,  the  rela- 
tive weights  of  a  molecule  of  chlorine  and  a  molecule  of  hydrogen 
chloride  stand  in  this  ratio.  The  molecule  of  chlorine  is  nearly  twice 
as  heavy  as  the  molecule  of  the  compound,  and  there  cannot  therefore 
be  a  whole  molecule  of  chlorine  in  a  molecule  of  hydrogen  chloride.  In 
fact,  we  perceive  at  once  that  the  molecule  of  hydrogen  chloride 


252 


INORGANIC   CHEMISTRY 


must  contain  only  half  a  molecule  of  chlorine  (35.46),  together  with 
half  a  molecule  of  hydrogen  (1).  In  other  words,  if  the  molecule  of 
free  chlorine  were  to  be  taken  as  the  atom  of  the  element,  then  the 
molecule  of  hydrogen  chloride  would  contain  only  half  an  atom  of 
chlorine,  which  would  be  contrary  to  our  decision  to  take  as  atoms 
quantities  which  are  not  divided.  So  we  choose  the  other  horn  of 
the  dilemma,  and  say  that  the  specimen  of  chlorine  in  the  molecule 
of  hydrogen  chloride  is  a  whole  atom  and  that  therefore  the  amount 
of  chlorine  in  the  molecule  of  free  chlorine  is  two  atoms,  and  its 
formula  C12.  Similarly,  the  weight  of  hydrogen  in  the  molecule  of 
hydrogen  chloride  is  1.008,  while  that  of  the  molecule  of  hydrogen 
is  2.016,  so  that  there  are  two  atoms  in  the  molecule  of  free  hydrogen 
and  its  formula  is  H2.  Reasoning  in  like  manner  from  the  molecular 
weights  of  oxygen  (32)  and  water  (18)  we  reach  the  conclusion  that 
the  molecule  of  oxygen  is  diatomic  (O2). 

Still  another  way  of  looking  at  the  same  facts  may  shed  light  on 
the  matter.  When  hydrogen  and  chlorine  combine,  one  volume  of 
each  of  these  gases  gives  two  volumes  of  hydrogen  chloride  (p.  228). 
Let  us  imagine  the  experiment  to  be  made  with  minute  volumes 
holding  one  hundred  molecules  each: 


HYDROGEN  CHLORIDE 


HYDROGEN        CHLORIDE 


100 

100 

came  from 


The  200  molecules  of  hydrogen  chloride  must  contain  at  least  200 
fragments  of  chlorine,  since  there  is  a  sample  in  each  molecule.  Now 
the  200  fragments  of  chlorine  came  from  a  volume  containing  only 
100  molecules  of  chlorine.  Each  of  the  latter  must  therefore  have 
been  split  in  the  chemical  action.  Hence  the  molecules  of  free 
chlorine  contain  at  least  two  atoms.  Parallel  reasoning  leads  to 
the  conclusion  that  the  molecules  of  free  hydrogen  are  likewise  di- 
atomic. If  we  consider  the  molecular  formula  of  a  substance  as 
representing  one  molecule  (see  below),  the  equation  for  this  action  is: 

H2  +  C12->2HC1. 

There  are  two  molecules  on  each  side  of  the  equation,  and  this  corre- 
sponds with  the  fact  that  there  is  no  change  in  the  total  volume. 
Again,  we  find  that  one  volume  of  oxygen  furnishes  enough  of  the 
element  for  two  volumes  of  water  vapor  (p.  156).  We  infer  there- 
fore that  each  molecule  of  oxygen  is  divided  into  two  parts  in  the 


MOLECULAR  WEIGHTS  AND  ATOMIC  WEIGHTS          253 

action.  And  in  like  manner,  when  we  find  that  one  volume  of  phos- 
phorus vapor,  in  combination  with  six  volumes  of  chlorine,  gives 
four  volumes  of  phosphorus  trichloride  vapor,  we  infer  that  every 
molecule  of  phosphorus  furnished  enough  of  the  element  for  four 
molecules  of  phosphorus  trichloride,  and  contained  therefore  four 
atoms  (see  Exercise  3,  p.  265). 


The  simple  fact  that  hydrogen  and  oxygen,  when  mixed,  do  not 
combine  (p.  126)  may  assist  in  reconciling  us  to  the  diatomic  nature 
of  their  molecules.  Some  part  of  the  mixture  has  to  be  heated 
strongly  to  start  the  interaction.  Now  the  molecular  formulae,  H2 
and  O2,  suggest  that  each  gas  is  really  in  combination  already  (with 
itself),  and  they  therefore  explain  to  some  extent  the  indifference  of 
the  gases  towards  one  another.  If  the  molecules  were  free  atoms, 
they  could  not  encounter  one  another  continually  as  they  move 
about,  and  yet  escape  combination  as  we  observe  that  they  do.  We 
may  imagine  that  the  primary  effect  of  heating  is  to  decompose 
some  of  the  molecules,  and  liberate  hydrogen  and  oxygen  in  the 
atomic  condition,  and  that  the  combination  of  these  atoms  starts 
the  explosion  of  the  whole  mass. 

In  the  case  of  hydrogen,  the  diatomic  nature  of  the  molecules 
has  been  demonstrated  by  an  entirely  different  method  by  Irving 
Langmuir.  It  has  long  been  known  that  the  conductivity  of  hydro- 
gen for  heat  is  greater  than  that  of  any  other  elementary  gas.  Thus, 
a  wire  raised  to  a  white  heat  in  air  by  means  of  an  electric  current 
cannot  be  kept  at  a  red  heat,  even,  by  the  same  current  in  hydrogen. 
In  other  gases,  heat  from  the  hot  wire  is  used  up  in  accelerating  the 
motion  of  the  molecules  of  the  gas.  Langmuir  has  shown,  however, 
that  in  hydrogen,  additional  heat  is  consumed  in  causing  decom- 
position of  many  of  the  diatomic  molecules  into  single  atoms: 
H2  <=^  2H.  He  has  measured  the  percentage  of  molecules  dissoci- 
ated (at  760  mm.),  and  found  that  it  varies  from  0.33  per  cent  at 
2000°  to  13  per  cent  at  3000°  and  34  per  cent  at  3500°.  When  the 
temperature  falls,  the  atoms  re-combine  to  form  diatomic  molecules. 

It  may  also  assist  in  making  the  matter  clear  if  we  note  that  the 
atomic  weight  of  an  element  is  the  unit  quantity  of  that  particular 
variety  of  matter,  when  it  is  in  combination.  The  unit  quantity  of 
the  same  variety  of  matter,  when  in  the  free  state,  as  a  substance, 
need  not  be  the  same.  We  should  not  expect  it  to  be  smaller,  but 
it  might  easily  be  twice  or  more  times  as  large. 


254  INORGANIC  CHEMISTRY 

Replies  to  Questions  about  Difficulties.  —  The  beginner 
always  becomes  confused  over  one  or  more  of  the  points  raised  by 
the  following  questions: 

1.  Why  was  32  g.  of  oxygen  taken  as  the  standard  for  molecular 
weights,  rather  than  16  g.?     Read  p.  252  and  footnote  to  p.  240. 

2.  If  02  is  the  smallest  mass  of  oxygen,  why  do  we  have  formulae 
like  H20  and  HC10?     02  is  the  smallest  mass  of  free  oxygen,  but  in 
combination  half  as  much  occurs  in  many  molecules.     Read  pp.  241, 
250,  and  251. 

3.  Why  is  not  the  atomic  weight  of  an  element  ascertained  by 
simply  measuring  the  density  of  the  elementary  substance?     Read 
pp.  251  and  261. 

4.  Can  we  not  deduce  the  valence  of  an  element  from  knowing 
the  number  of  atoms  in  its  molecule,  and  vice  versa?     Some  molec- 
ular formulae  and  valences  are:    IV,  O2n,  CV,  Zn11,  also  Hg  (uni- 
valent  and  bivalent) ,  ?4  (trivalent  and  quinquivalent)  and  Sg  (bivalent 
and  sexivalent).     There  is  no  relation,  either  observable  or  to  be 
expected. 

5.  Do  the  molecular  weights,  oxygen  =  32  and  hydrogen  =  2, 
mean  that  the  molecules  of  oxygen  are  larger  than  are  those  of 
hydrogen?     This  is  the  ratio  of  their  weights,  but  none  of  the  phe- 
nomena discussed  in  this  chapter  are  influenced  appreciably  by  their 
relative  sizes,  and  therefore  none  of  them  give  any  information  on 
the  subject.     Read  the  footnote  to  p.  237. 

Exercises.  —  1.  The  weight  of  1  1.  of  gas  at  0°  and  760  mm.  is 
5.236  g.  What  is  the  density  referred  (a)  to  air  (air  =  1)  and  (6)  to 
hydrogen,  and  (c)  what  is  the  molecular  weight  (pp.  233,  237)? 

2.  The  density  of  a  gas,  referred  to  air,  is  6.7.     What  is  the 
weight  of  1  1.  (p.  233),  and  what  is  the  molecular  weight  (p.  237)? 

3.  The  molecular  weight  of  a  substance  is  65.     What  is  the 
density  referred  to  air,  and  what  is  the  weight  of  1  1.? 

4.  If  the  molecular  weight  of  oxygen  were  taken  as  100,  what 
would  be  the  volume  of  the  G.M.V.  (p.  233)?    What,  on  the  same 
scale,  would  be  the  molecular  weight  of  water,  and  what  would  be 
the  atomic  weights  of  hydrogen  and  chlorine  (pp.  233,  241)? 


CHAPTER  XIII 

APPLICATIONS   OF  MOLECULAR  AND  ATOMIC   WEIGHTS. 
PROPERTIES   OF  ATOMS 

Applications:  Interactions  Between  Gases.  —  According  to 
Avogadro's  law,  if  we  filled  a  succession  of  vessels  of  equal  dimen- 
sions with  different  gases,  and  could  arrest  the  motion  of  the  particles 
and  observe  their  disposition,  we  should  find  that  the  average  dis- 
tance from  particle  to  particle  would  be  the  same  in  all  cases.  This 
would  be  true  whether  our  vessels  were  filled  with  single  gases,  with 
homogeneous  mixtures,  or  with  gases  in  layers.  Such  being  the  case, 
if  any  chemical  change  is  brought  about  in  the  mass  which  results 
in  a  multiplication  of  the  molecules,  it  is  evident  that  the  volume 
will  have  to  increase  in  order  that  the  spacing  may  remain  the  same 
as  before.  If  any  chemical  action  results  in  a  diminution  of  the 
number  of  molecules,  then  a  shrinkage  must  take  place  in  order  that 
the  spacing  may  be  preserved  as  before.  Thus,  in  a  mixture  of 
hydrogen  and  chlorine,  according  to  our  hypothesis,  when  interaction 
to  produce  hydrogen  chloride  occurs,  neighboring  molecules  of  hy- 
drogen and  chlorine  simply  exchange  units,  so  that  HH  -f-  C1C1 
becomes  HC1  ~f  C1H.  There  being  no  alteration  in  the  number  of 
particles,  no  change  in  volume  occurs.  In  the  case  of  water,  on  the 
other  hand, 

HH  +  OO  +  HH    becomes    HOH  +  HOH. 

Since  the  oxygen  molecules,  which  form  a  third  of  the  whole,  dis- 
appear into  the  molecules  of  hydrogen,  the  tendency  to  preserve 
spacing  results  in  a  diminution  of  the  volume  by  one-third  (p.  156). 
Thus  Gay-Lussac's  law  would  have  followed  as  a  natural  inference 
from  Avogadro's  law,  if  the  former,  being  more  obvious,  had  not 
been  discovered  first. 

This  method  of  looking  upon  chemical  interactions  between 
gases  gives  us  the  nearest  sight  which  we  can  have  of  the  behavior  of 
the  molecules  themselves.  We  cannot  perceive  the  individual  mole- 
cules, but,  in  consequence  of  the  spatial  arrangement  which  they 
observe,  the  change  in  the  whole  volume  of  a  large  aggregate  of 

255 


256  INORGANIC  CHEMISTRY 

molecules  enables  us  to  draw  conclusions  at  once  in  regard  to  the 
behavior  of  the  single  molecules  in  detail. 

Applications:  Molecular  Equations.  —  To  utilize  the  fore- 
going considerations,  chemists  always  employ  in  their  equations  the 
molecular  formulae  for  the  gases  and  easily  vaporized  substances 
concerned.  Thus  far,  we  have  used  the  equation: 

H  +  Cl  ->  HC1, 
WEIGHTS:  1.008    35.46       36.468 

and  the  information  it  contained  was  exhausted  when  we  had  placed 
below  the  symbols  the  weights  for  which  they  stood.  But  the  molec- 
ular equation  is  much  more  instructive.  The  following  shows  the 
interpretations  to  which  the  molecular  equation  is  subject: 

H2      +      C12    -»      2HC1. 

WEIGHTS:  2.016  g.         70.92  g.         2  X  36.468  ( =  72.936)  g. 

VOLUMES:  22.4 1.  22.4 1.  2  X  22.4 1. 

MOLECULES:  11  2 

'"J.  • 
The  weights,  although  doubled,  show  the  same  proportions,  so 

that  questions  of  weight  are  answered  as  easily  as  before.  These 
weights,  however,  being  molecular  weights,  or  multiples  thereof,  can 
be  translated  at  once  into  volumes,  and  questions  about  volumes 
can  also  be  answered.  Finally  the  relative  numbers  of  each  kind  of 
molecules  can  be  read  from  this  equation,  for  the  coefficients  in  front 
of  the  formulae  represent  these  numbers.  Where  no  coefficient  is 
written,  1  is  to  be  understood.  The  application  of  these  properties 
of  molecular  equations  is  illustrated  below.  Before  applying  these 
equations,  however,  we  must  first  learn  how  to  make  them. 

Applications:    The  Making  of  Molecular  Equations.  —  To 

make  a  molecular  equation,  we  first  make  an  equation  according  to 
the  rules  already  explained  (p.  73).  An  equation  like  that  given  for 
the  interaction  of  oxygen  with  hydrogen  chloride  (Deacon's  process, 
p.  216):  2HC1  +  O  ->  H2O  +  2C1,  is  the  result.  Then  we  adjust 
the  equation  so  that  molecular  formulae  are  used  throughout.  2CI 
becomes  at  once  C12.  The  oxygen,  however,  must  also  appear  as  O2, 
or  a  multiple  of  this,  in  such  equations.  Hence  the  whole  equation 
must  be  multiplied  by  2: 

4HC1  +•  O2  -»  2H2O  +  2C12. 


APPLICATIONS  257 

Again,  the  equation  for  the  preparation  of  chlorine  from  potassium 
permanganate  and  hydrochloric  acid  (p.  218)  becomes: 

2KMn04  +  16HC1  ->  8H2O  +  2KC1  +  2MnCl2  +  5C12. 

Every  equation  containing  an  odd  number  of  atoms  of  a  substance 
whose  molecules  are  diatomic  must  be  multiplied  by  2.  Again,  mer- 
curic oxide  decomposes  to  give  mercury  vapor  and  oxygen  (p.  17), 
and  the  molecules  of  mercury  are  monatomic  and  those  of  oxygen 
diatomic,  so  we  write: 


Finally  the  formulae  of  substances  which  are  solid  or  liquid,  and 
cannot  be  easily  vaporized,  are  written  in  the  simplest  terms.  Thus, 
since  substances  like  the  copper  in  the  following  equation  are  in- 
volatile,  the  molecular  weights  of  such  substances  are  unknown,  and 
their  molecular  formulae  likewise:  2Cu  +  02  —  >  2CuO.  Further- 
more, in  the  case  of  substances  which  can  be  volatilized,  although 
the  molecular  weights  and  molecular  formulae  may  therefore  be 
known,  we  do  not  usually  employ  the  molecular  formulae  if  the  sub- 
stance is  not  used  in  the  form  of  vapor  in  the  laboratory.  Thus,  the 
molecular  formula  of  phosphoric  anhydride  is  P4Oio  (p.  88).  But 
we  generally  make,  and  use,  only  the  solid  form,  and  not  the  vapor, 
in  actual  work.  Hence  the  action  with  water  is  usually  written  as 
we  have  given  it  (p.  149),  rather  than  in  the  form:  P^io  +  6H«jO 
rn*  4H3P04. 

The  applications  of  the  properties  of  molecular  equations  (which 
will  be  used  exclusively  hereafter)  may  now  be  illustrated  in  detail. 

Applications:  To  Arithmetical  Problems.  —  1.  When  a 
problem  in  regard  to  weights  of  material  used  or  produced  in  a  given 
action  is  to  be  solved,  the  molecular  equation  is  to  be  written  and 
the  weights  inserted  beneath  the  formulae.  The  mode  of  calculation 
has  been  described  already  (p.  75). 

2.  When  a  problem  involving  weights  and  volumes  is  to  be 
solved,  the  molecular  equation  is  to  be  written,  and  both  the  weights 
and  volumes  are  to  be  inserted.  Note,  however,  that  the  volumes 
of  the  substances  which  are  in  the  gaseous  condition  only  are 
inserted. 

For  example:  What  volume  of  oxygen  is  obtained  from  60  g.  of 
potassium  chlorate?  The  molecular  equation,  made  as  shown  above 
(p.  256),  together  with  the  full  interpretation,  are  as  follows: 


258  INORGANIC  CHEMISTRY 


WEIGHTS: 


2KC1O3        ->  2KC1      +      302. 

2  (39.1  +  35.46  +  48)       2  (39.1  +  35.46)          3  X  32 


245.12  g.  149.12  g.  96  g. 

VOLUMES:  3X22.41. 

GIVEN:  60  g.  x  1. 

Observe  that  no  volumes  are  given  under  the  chlorate  and  chloride 
of  potassium.  This  is  because  their  volumes  in  the  gaseous  condition 
can  be  of  no  practical  use,  since  they  are  solids  which  are  melted,  but 
not  vaporized  during  this,  or  any  action  in  which  we  employ  them. 
Note  also  that  the  data  given  in  the  problem  are  inserted,  and  x  1. 
is  placed  under  the  oxygen.  Now,  as  to  the  problem  in  hand,  it  is 
concerned  with  a  weight  of  potassium  chlorate  and  a  volume  of 
oxygen.  Reading  from  the  equation,  our  information  on  these  points 
is  that  245.12  g.  of  potassium  chlorate  give  67.2  liters  (observe  that 
the  coefficients  are  used,  as  well  as  the  molecular  weights,  in  these 
numbers)  of  oxygen  at  0°  and  760  mm.,  and  the  question  is,  What 
volume  will  60  g.  give?  By  proportion,  245.12  g.  :  67.2  1.  :  :  60  g.  :  x 
1.,  where  x  =  16.4  liters.  If  a  different  temperature  and  pressure 
had  been  specified,  either  the  volume  in  the  equation,  or  the  answer, 
would  have  had  to  be  converted  by  rule  to  the  given  conditions. 

It  saves  time  not  to  write  out,  as  above,  the  whole  interpretation, 
but  only  the  parts  required.  For  example,  if  the  question  is  :  What 
volume  of  chlorine  is  needed  to  give  25  g.  of  aluminium  chloride,  we 
may,  if  we  choose,  omit  all  the  data  excepting  the  volume  of  the 
chlorine  and  the  weight  of  the  aluminium  chloride,  thus: 

2A1  +     3C12     —  >     2AlCla 

3  X  22.4  1.     2  X  133.48  g. 
xl  25  g. 

The  volume  of  chlorine  required  is  3  X  22.4  X  25  -f-  (2  X  133.48) 
liters.  These  illustrations  show  the  method  of  calculating  actual 
volumes  (see  Exercises  4,  5). 

3.  If  the  question  concerns  relative  volumes  only,  then  it  is 
simplest  to  use  the  interpretation  of  the  equation  in  terms  of  mole- 
cules. For  example,  What  relative  volumes  of  hydrogen  chloride 
and  oxygen  are  required  in  Deacon's  process?  The  molecular  equa- 
tion is  (p.  256)  : 


MOLECULES: 


APPLICATIONS  259 

Since  equal  numbers  of  molecules  of  gases  occupy  equal  volumes,  the 
proportion  4  molecules  of  hydrogen  chloride  to  1  molecule  of  oxygen 
shows  the  ratio  to  be  4  :  1  by  volume.  Similarly,  every  4  molecules 
of  hydrogen  chloride  give  2  molecules  of  chlorine,  so  that  the  ratio  of 
these  substances  by  volume  is  4  :  2,  or  2  :  1. 

In  regard  to  the  water,  since  that  is  not  a  gas  at  common  tempera- 
tures, the  question,  if  asked,  must  be  more  specific:  What  are  the 
relative  volumes  of  steam  and  chlorine  in  the  product,  as  commonly 
delivered  by  the  action  at  400°?  It  is  2  :  2,  or  1  :  1.  What  are  the 
relative  volumes  of  water  and  chlorine,  after  the  products  have  cooled 
to  room  temperature?  The  water  is  no  longer  a  gas,  so  that  it  occu- 
pies, relatively,  almost  no  volume.* 

What  is  the  total  volume-change  in  the  foregoing  action  above 
100°?  It  is  a  change  from  5  molecules  to  4.  The  volume,  changes  in 
the  same  ratio.  But  at  0°  the  volume-change  is  from  5  volumes  to  2, 
for  the  water  does  not  appreciably  add  to  the  volume  of  the  products 
(see  Exercises  6,  7). 

4.  When  we  know  the  molecular  formulae  of  the  single  substances 
concerned  in  an  action,  the  equation  can  be  made,  and  the  relative 
volumes  determined,  without  actual  measurement.  For  example: 
What  volume-change  will  be  observed  when  a  mixture  of  carbon 
monoxide  and  oxygen  has  exploded,  and  the  temperature  has  once 
more  reached  that  of  the  room?  The  molecular  formulae  are  CO,  O2, 
and  CO2.  The  equation  representing  the  weights  is  CO  +  O  —  >  C02. 
The  molecule  of  oxygen,  however,  being  O2,  we  cannot  employ  less 
than  this  quantity  in  a  molecular  equation,  so  that  the  equation 
becomes  : 


Three  molecules,  therefore,  give  two,  throughout  the  whole  mass,  and 
therefore  three  volumes  will  become  two,  if  the  pressure  and  tempera- 
ture are  the  same  at  the  beginning  and  end  of  the  action. 

If  we  remember  that  all  volatile  compounds  of  carbon  and  hydro- 
gen burn  to  .form  water  and  carbon  dioxide,  the  molecular  equation 

*  Of  course  if  an  exact  answer  must  be  given,  it  can  be  given.  But  for  this 
we  require  the  weight  and  specific  gravity  of  the  product.  Thus,  2H2O  represents 
2  X  18  g.  of  water.  The  sp.  gr.  of  water  is  1.  Therefore  the  volume  of  water 
formed  is  36  c.c.  The  volume  of  2C12  is  2  X  22.4,  or  44.8  liters  at  0°.  The  ratio 
of  water  to  chlorine  by  volume  at  0°  is  therefore  36  :  44,800.  But,  as  a  rule,  we 
simply  give  the  volumes  of  solids  and  liquids  as  zero,  compared  with  those  of  the 
gases  concerned  in  the  same  action. 


260  INORGANIC  CHEMISTRY 

for  any  such  combustion  may  easily  be  made,  and  the  volumes  of  all 
the  materials  ascertained.  When  water  is  a  product,  only  its  volume 
as  steam  is  given  by  the  equation  (see  Exercises  7,  8). 

5.  Knowing  by  heart  the  molecular  formulae  of  gaseous  sub- 
stances, as  we  must  know  them  for  many  purposes,  it  is  unnecessary 
to  burden  our  minds  with  other  data  in  regard  to  the  relative  den- 
sities of  gases.  Is  hydrogen  chloride  HC1  heavier  or  lighter  than 
carbon  dioxide  CO2?  These  formulas  represent  the  weights  of  equal 
volumes  (22.4  1.),  namely  36.46  g.  and  44  g.,  respectively.  Hence 
the  former  gas  is  a  little  lighter.  Remembering  that  the  G.M.V.  of 
air  weighs  28.955  g.  (Table,  p.  233),  we  can  compare  the  weight  of 
any  gas  with  that  of  air  in  the  same  way. 

What  are  the  relative  weights  of  acetylene  (C2H2,  p.  241)  and 
sulphur  dioxide  S02  as  compared  with  air?  The  G.M.V.  cube  holds 
formula-weights  of  the  first  two,  namely  26  g.  and  64  g.,  and  28.955  g. 
of  air.  Hence  acetylene  is  a  little  lighter  than  air,  and  sulphur 
dioxide  more  than  twice  as  heavy  (see  Exercise  9). 

Applications:  To  Cases  of  Dissociation.  —  Several  gases  or 
vapors  yield  smaller  values  for  their  densities,  and  therefore  molec- 
ular weights,  when  the  densities  are  measured  at  higher  tempera- 
tures. This  indicates  that  the  molecules  have  become  lighter,  and 
can  only  mean  that  decomposition  has  taken  place  in  consequence 
of  the  heating.  Behavior  of  this  kind  is  shown  both  by  compounds 
and  by  simple  substances. 

For  example,  phosphorus  pentachloride  PC15,  although  a  solid, 
can  be  converted  into  vapor  without  much  difficulty.  Its  molecular 
weight,  if  it  underwent  no  chemical  change  during  the  volatilization, 
would  be  31  +  177.3  =  208.3.  The  density  actually  observed  at 
300°  and  760  mm.  pressure  gives  by  calculation  not  much  more  than 
half  this  value.  The  direct  inference  from  this  is,  that  the  molecules 
have  only  half  the  (average)  weight  that  we  expected:  or,  in  other 
words,  are  twice  as  numerous  as  we  expected.  The  explanation  is 
found  when  we  examine  the  nature  of  the  vapor  more  closely.  We 
find  that  it  is  a  mixture  of  phosphorus  trichloride  and  free  chlorine, 
resulting  from  a  chemical  change  according  to  the  equation :  PCls  ^ 
PC13  +  C12.  The  low  value  of  the  density  thus  tells  us  that  disso- 
ciation has  taken  place.  From  the  value  of  the  density  at  various 
temperatures,  we  may  even  calculate  the  proportion  of  the  whole 
material  which  is  dissociated.  At  300°  it  is  97  per  cent;  at  250°,  80 
per  cent;  and  at  200°,  48.5  per  cent.  Thus,  when  the  temperature 


APPLICATIONS  261 

is  lowered,  progressive  recombination  takes  place  and  the  proportion 
dissociated  becomes  less.  Finally  the  vapor  condenses  and  yields 
the  original  solid. 

Again,  sulphur  boils  at  445°,  but  is  easily  vaporized  at  a  tempera- 
ture as  low  as  193°,  under  very  low  pressure.  At  this  temperature 
the  density  of  the  vapor  gives  the  molecular  weight  256  (=  8  X  32), 
and  the  molecular  formula  S8.  That  is  to  say,  the  G.M.V.  holds 
256  g.  of  the  vapor  at  193°.  At  800°,  however,  the  density  is  only 
one-fourth  as  great,  and  the  G.M.V.  holds  only  64  g.  (S2).  This 
means  that  256  g.  now  occupy  four  times  as  large  a  volume  as  before, 
and  the  increase  is  additional  to  the  effect  of  the  mere  thermal  expansion, 
which  is  allowed  for  in  the  calculation  and  eliminated.  Hence  the 
molecules  have  dissociated.  At  1700°  the  molecular  formula  is  still 
S2,  so  that  this  shows  the  limit  of  observed  dissociation :  Ss  <=*  4S2. 
When  the  vapor  is  cooled,  the  density  increases  once  more  and  at  193° 
recovers  completely  the  greater  value.  Similar  observations  show 
that  phosphorus  vapor  at  313°  is  all  P4,  but  at  1700°  one-half 
of  the  molecules  are  P2.  Iodine  vapor,  up  to  700°,  is  all  I2.  Beyond 
this  temperature  the  density  diminishes,  and  when  1700°  is  reached 
the  vapor  is  all  I.  Thus  the  molecules  are  diatomic  at  low  tempera- 
tures and  monatomic  at  high  ones.  The  densities  of  oxygen,  hydro- 
gen, and  chlorine  are  not  measurably  affected  by  heating  to  1700°, 
so  that  their  diatomic  molecules  exist  from  temperatures  far  below 
0°  up  to  1700°,  and  are  evidently  very  stable.  For  observations  on 
hydrogen  above  1700°,  however,  see  p.  253. 

Applications:  Finding  the  Atomic  Weight  of  a  New  Ele- 
ment. —  By  way  of  reviewing  the  principles  explained  in  this  chap- 
ter, let  us  apply  them  to  the  imaginary  case  of  a  newly  discovered 
element.  The  bromide  of  the  element  is  found  to  be  easy  of  prepa- 
ration and  to  be  volatile.  The  bromide  contains  30  per  cent  of  the 
element,  and  therefore  70  per  cent  of  bromine,  and  its  vapor  density 
referred  to  air  is  11.8.  The  analysis  can  always  be  made  much  more 
accurately  than  the  measurement  of  vapor  density,  so  that  the  two 
former  numbers  are  more  trustworthy  than  the  last. 

To  find  the  equivalent  of  the  element,  that  is,  the  amount  com- 
bined with  79.92  parts  (the  equivalent,  and  also  the  atomic  weight), 
of  bromine,  we  have  the  proportion  70  :  30  :  :  79.92  :  x,  from  which 
x  =  34.3.  The  atomic  weight  must  be  this,  or  some  small  multiple 
of  it. 

The  G.M.V.  of  air  weighs  28.955  g.  (p.  233).     Hence  the  same 


262  INORGANIC  CHEMISTRY 

volume  of  the  vapor  of  this  bromide,  which  is  11.8  times  as  heavy  as 
air,  will  weigh  28.955  X  11.8,  or  341.67  g.  This  is  therefore  the 
molar  weight  of  the  compound. 

Now  30  per  cent  of  this  is  the  new  element:  341.67  X  30  -f-  100 
=  102.5.  Three  times  the  equivalent  weight  is  the  multiple  nearest 
to  this  number,  3  X  34.3  =  102.9,  the  difference  being  due  to  error  in 
determining  the  density.  So  long  as  no  other  volatile  compound  is 
known,  we  adopt  this  as  the  atomic  weight.  The  rest  of  the  molar 
weight  (239.76  =  3  X  79.92  parts)  is  bromine.  Thus  the  formula  of 
the  compound  is  ElBr3,  and  from  this  we  see  that  the  element  is 
trivalent. 

In  case  no  volatile  compound  of  the  element  can  be  formed,  the 
weight  combining  with  79.92  parts  of  bromine  is  measured  as  before. 
Then  some  of  the  free  simple  substance  is  made,  say  by  electrolysis, 
and  its  specific  heat  is  determined.  The  sp.  ht.  is  about  0.063. 
Application  of  Dulong  and  Petit's  law  then  gives  the  atomic  weight. 
The  product  34.3  X  0.063  is  equal  to  2.161.  Hence,  the  equivalent 
must  be  multiplied  by  3  to  give  the  atomic  weight,  for  this  raises 
the  product  to  6.48,  which  is  within  the  limits.  Thus  the  value  of 
the  atomic  weight  is  102.9,  as  before. 

PROPERTIES  OF  ATOMS 

There  is  danger  that  one  may  attribute  to  atoms  properties 
which  they  do  not  possess.  It  is  sometimes  said  that  an  atom  of  an 
element  is  the  smallest  particle  of  that  element  which  possesses  the 
same  properties  as  a  larger  portion  of  the  same  element.  This  is  an 
ill-considered  statement.  For  example,  an  atom  of  mercury,  in 
mercury  vapor,  has  the  perfect  elasticity  of  a  molecule,  but  a  mass 
of  frozen  mercury  has  a  limited  degree  of  elasticity,  like  any  other 
piece  of  matter.  We  should  note,  therefore,  the  properties  which 
the  experimental  work  thus  far  described  has  shown  atoms  to  possess. 

Properties  of  Atoms.  —  Atoms  move  from  one  state  of  combi- 
nation to  another  without  alteration  in  their  mass.  Since  they  may 
be  restored  to  their  free  condition  and  recombined  as  often  as  we 
choose,  without  impairment  of  their  individuality,  each  kind  must  be 
composed  of  a  distinct  variety  of  matter. 

The  conception  of  valence  (p.  130)  suggests  that  some  atoms 
unite  with  but  one  other  atom,  habitually  (NaCl).  Some,  however, 
unite  with  two  of  the  first  kind  (ZnCl2),  or  with  one  other  of  their 


PROPERTIES  OF  ATOMS  263 

own  kind  (ZnO),  still  others  with  three  atoms  of  the  first  kind 
(A1C13),  and  so  forth.  In  other  words,  it  involves  the  assumption 
that  each  kind  of  atom  has  a  limited  capacity  for  holding  other 
atoms  in  combination.  Thus,  taking  the  most  crudely  mechanical 
view  of  the  matter,  we  might  suggest  that  there  is  a  limit  to  the 
number  of  points  at  which  atoms  may  be  attached  to  one  another. 
When  one  atom  of  chlorine  is  attached  to  one  of  sodium,  the  com- 
bining capacity  of  each  is  exhausted.  When  one  atom  of  hydrogen 
is  attached  to  one  atom  of  oxygen,  one  combining  capacity  of  the 
oxygen  still  remains  (H  —  O  — ),  and  can  be  satisfied  by  one  more 
atom  of  hydrogen  or  of  some  other  element. 

Atoms  of  the  same  kind  can  combine  with  one  another.  Thus 
we  have  H2,  O2,  Os,  N2,  and  so  forth.  But  many  kinds  of  atoms 
lack  this  quality.  The  molecules  of  metals,  so  far  as  we  have  been 
able  to  ascertain  their  nature,  are  monatomic,  e.g.,  K,  Na,  Hg,  Zn,  Cd. 
Then  some  atoms,  those  of  the  noble  gases  (q.v.),  refuse  to  combine 
with  any  other  atoms  whatever,  e.g.,  He,  Ne,  A,  Kr,  Xe,  Nt. 

Again,  certain  compounds,  similar  in  composition,  such  as  barium 
peroxide  BaO2  and  manganese  dioxide  MnO2,  show  different  chem- 
ical behavior.  The  former  of  these  when  treated  with  sulphuric 
acid  gives  hydrogen  peroxide  (q.v.),  while  the  latter  gives  only 
water  and  oxygen.  If  it  were  merely  a  question  of  composition, 
the  substances  should  behave  alike.  So,  also,  we  often  have  two  or 
more  compounds  identical  in  molecular  weight,  and  in  the  elements 
and  numbers  of  units  which  they  contain,  which  are  nevertheless 
totally  different  in  physical  and  chemical  properties.  To  account 
for  this  difference  we  have  found  it  convenient  to  suppose  that, 
although  the  atoms  contained  in  the  two  or  more  kinds  of  molecules 
are  of  the  same  numbers  and  kinds,  the  atoms  are  in  different  geo- 
metrical arrangement  towards  one  another.  Sometimes  one  of  these 
substances  can  be  made  directly  from  the  other,  and  this  gives  us 
that  variety  of  chemical  change  which  was  named  internal  rear- 
rangement (p.  20).  Thus,  ammonium  cyanate  NH4CNO,  when 
heated  in  aqueous  solution,  turns  into  urea  CO(NH2)2,  a  substance 
with  totally  different  chemical  properties. 

Finally,  the  atoms  of  uranium,  thorium,  radium,  and  some  other 
elements,  disintegrate  spontaneously,  giving  off  atoms  of  helium  and 
electrons  (see  Radioactivity),  so  that  some  kinds  of  atoms  are 
unstable. 

To  sum  up:  Each  kind  of  atom  has  a  specific  mass,  and  consists 
as  a  whole  of  a  specific  kind  of  material.  Each  atom,  of  the  kinds 


264  INORGANIC  CHEMISTRY 

that  can  enter  into  combination,  has  a  maximum  capacity  for  hold- 
ing other  atoms  (valence).  When  there  are  more  than  two  atoms  in 
the  molecule  of  a  compound,  there  is  often  more  than  one  way  of 
connecting  them  so  as  to  satisfy  the  valence,  and  thus  two  or  more 
different  substances  of  the  same  composition  may  exist.  Some 
atoms  are  unstable,  and  we  know  something  of  the  structure  of  all 
atoms  (see  Chap.  XXII). 

Properties  of  Compounds  not  Explained  by  Properties  of 
Atoms.  —  In  one  direction,  namely,  that  of  accounting  for  the 
properties  of  compounds,  no  attempt  has  been  made  to  adapt  the 
atomic  viewpoint  so  as  to  explain  the  facts.  Thus,  two  hydrogen 
atoms  in  a  molecule  give  a  gas  almost  insoluble  in  water;  two  chlo- 
rine atoms,  a  gas  which  is  moderately  soluble;  but  one  of  each  gives 
hydrogen  chloride,  which  dissolves  in  water  in  extraordinary  quan- 
tities. So  also,  colorless  substances  give,  by  chemical  union,  strongly 
colored  ones,  and  odorless  substances,  by  chemical  union,  strongly 
odorous  ones.  Putting  pieces  of  iron  and  sulphur  side  by  side 
causes  absolutely  no  change  in  the  properties  of  either.  And  yet 
the  atomic  viewpoint  compels  us  to  assume  that  if  the  particles  are 
made  fine  enough,  and  placed  close  enough  to  one  another,  the 
individual  properties  of  the  constituents  will  entirely  disappear. 
Hitherto  we  have  failed  to  think  of  any  qualities  which  might  be 
attributed  to  the  atoms  in  order  to  account  for  facts  of  this  class. 
Why  should  oxygen  O2  and  ozone  O3  be  so  different  in  behavior 
although  the  atomic  viewpoint  hints  at  nothing  but  a  substitution 
of  three  atoms  for  two?  What  atomic  properties  shall  account  for 
the  difference  between  red  and  yellow  phosphorus? 

Definitions  Should  be  in  Experimental  Terms.  —  Defini- 
tions are  often  stated  in  terms  of  atoms  and  molecules.  For  ex- 
ample: A  physical  change  is  one  in  which  the  relations  amongst  the 
molecules  are  altered,  but  the  molecules  themselves  are  not  affected 
in  any  way;  while  a  chemical  change  is  one  in  which  the  molecules 
themselves  are  decomposed,  or  their  contents  are  rearranged.  State 
this  definition  to  some  one,  and  instruct  him  then  to  examine  a  vessel 
of  boiling  water,  with  the  definition  in  mind,  and  report  whether 
the  process  of  vaporization  is  physical  or  chemical.  He  will  report 
that  he  could  not  perceive  either  molecules  or  atoms  in  the  vessel, 
and  therefore  cannot  answer  the  question.  Definitions  like  the 
above  can  be  memorized  and,  with  the  help  of  the  imagination,  they 


PROPERTIES  OF  ATOMS  265 

can  be  understood,  but  they  are  not  directly  applicable  to  any 
question  arising  in  the  laboratory. 

Again,  the  law  of  definite  proportions:  Each  species  of  molecules 
contains  the  same  kind  and  number  of  atoms.  We  cannot  directly 
count  and  classify  the  atoms  in  ten  specimens  of  potassium  chlorate, 
to  ascertain  whether,  in  this  respect,  they  are  all  alike.  Such  defi- 
nitions require  elaborate  translation  into  other  terms  before  they 
can  be  connected  with  any  practicable  method  of  verifying,  illus- 
trating, or  employing  them.  Chemistry  is  a  practical  subject,  and 
the  statement  and  definitions  should  all  be  in  terms  of  experimental 
facts  and  phenomena,  and  should  never  require  translation  into 
other  terms  before  they  can  be  put  to  practical  use.  Hence  the 
definitions  we  have  given  (cf.  p.  54)  have  all  been  in  terms  of  experi- 
ments. 

Exercises.  —  1.  The  chloride  of  a  new  element  contains  38.11 
per  cent  of  chlorine  and  61.89  per  cent  of  the  element.  The  vapor 
density  of  the  compound  referred  to  air  is  12.85.  What  is  the  atomic 
weight  of  the  element,  so  far  as  investigation  of  this  one  substance 
can  give  it  (p.  261)?  What  is  its  valence? 

2.  In  future  nothing  but  molecular  formula  of  free  elements  must 
be  used  (p.  250).     Write  in  molecular  form  ten  of  the  equations  in- 
volving gases  which  are  found  in  the  preceding  chapters. 

3.  If  a  new  form  of  oxygen  were  found,  such  that  one  volume  of 
it  required  four  volumes  of  hydrogen  to  produce  water,  what  would 
be  its  molecular  formula  (p.  252)?     What  would  be  the  weight  of 
22.4  1.? 

4.  What  volume  of  oxygen  at  10°  and  750  mm.  is  obtainable  by 
heating  50  g.  of  barium  peroxide  (pp.  82,  257)? 

5.  What  volume  of  oxygen  at  20°  and  760  mm.  is  required  to  con- 
vert 16  g.  of  iron  into  dehydrated  rust  Fe2O3  (p.  257)? 

6.  Write  out  the  molecular  equations  for  the  interactions  of 
methane  and  chlorine  giving  CH3C1;    and  for  the  burning  of  phos- 
phorus (vapor)  in  oxygen  (p.  257).     Deduce  the  volume  relations  of 
the  initial  substances,  and  of  the  products,  at  0°  and  100°  in  each 
case. 

7.  Write  out  the  molecular  equations  for  the  interactions  of 
acetylene  and  oxygen  (p.  259),  and  of  alcohol  vapor  (b.-p.  78°)  and 
oxygen.     Deduce  the  volume  relations  of  the  initial  substances  and 
of  the  products  at  0°  and  at  100°  in  each  case. 

8.  The  molecular  weight  of  cyanogen  is  52.08.     What  is  its  den- 


266  INORGANIC  CHEMISTRY 

sity  referred  to  air,  and  what  the  weight  of  1  1.  at  0°  and  760  mm.? 
It  contains  46.08  per  cent  carbon  and  53.92  per  cent  nitrogen.  What 
is  the  formula  of  the  substance  (p.  72)?  Exploded  with  oxygen  it 
forms  carbon  dioxide  and  free  nitrogen.  What  will  be  the  relative 
volumes  of  the  materials  before  and  after  the  interaction  (p.  259)? 

9.  What  are  the  relative  weights  of  equal  volumes  of  hydrogen 
sulphide  H2S,  and  hydrogen  iodide  HI,  compared  with  air  (p.  258)? 

10.  At  1700°  the  average  molecular  weight  of  phosphorus  is  91 
(p.  261).     What  percentage  of  molecules  of  ?4  has  been  dissociated 
into  P2? 

11.  Show  that,  if  an  element  has  more  than  one  equivalent  weight, 
the  atomic  weight  must  be  some  multiple  of  each  of  the  equivalents 
by  a  whole  number. 

12.  Prove  that  16  is  preferable  to  8  for  the  atomic  weight  of 
oxygen,  because  the  smaller  number  involves  a  fractional  value  for 
the  atomic  weight  of  hydrogen. 

13.  In  previous  chapters  our  definitions  have  been  experimental. 
In  imitation  of  the  definitions  of  the  law  of  definite  proportions  and 
of  valence  (p.  264),  give  theoretical  definitions  of  the  following,  in 
terms  of  the  atomic  hypothesis :   (a)  Multiple  proportions,  (b)  chem- 
ical unit  weight,  (c)  molecular  weight,  (d)  element,  (e)  compound, 
(/)  symbol,  (g)  formula,  (h)  equation. 

14.  Criticize  the  definitions:    (a)  The  atomic  weight  of  an  ele- 
ment is  the  smallest  portion  of  that  element  which  takes  part  in 
chemical  change,     (b)  An  atom  is  the  smallest  particle  that  can  be 
conceived. 

15.  Define  all  the  varieties  of  chemical  change  (p.  21)  in  terms 
of  the  atomic  hypothesis. 


CHAPTER  XIV 

THE   HALOGEN  FAMILY 
FLUORINE,   CHLORINE,   BROMINE,   IODINE 

THE  elements  to  which  we  have  so  far  devoted  most  attention  have 
been  oxygen,  hydrogen,  and  chlorine.  If  we  recall  the  chemical  prop- 
erties and  relations  of  these  elements  we  shall  recognize  the  fact  that 
they  all  possess  very  distinct  individualities. 

The  Chemical  Relations  of  Elements.  —  Hydrogen  is  the 
substance  (p.  127)  which  unites  readily  with  oxygen  and  chlorine,  less 
readily  with  other  non-metals,  and  scarcely  at  all  with  metals.  Oxy- 
gen and  chlorine  resemble  one  another  somewhat  in  the  greatness  of 
their  chemical  activity  and  the  variety  of  free  elements  with  which 
they  are  capable  of  uniting,  but  differ  markedly  in  what  we  have 
called  their  chemical  relations  (p.  226).  The  resulting  compounds 
belong,  in  fact,  to  quite  different  classes  —  oxygen  forms  oxides,  chlo- 
rine forms  chlorides  —  and  elements  are  considered  similar  only  when 
they  resemble  one  another  in  chemical  relations,  and  produce,  by  com- 
bination with  the  same  element,  compounds  having  similar  chemical 
properties.  Thus  the  common  oxide  of  hydrogen,  water,  is  a  neutral 
substance,  and  is  chemically  rather  indifferent.  The  chloride  of 
hydrogen  in  aqueous  solution  is  a  strong  acid  and  is  chemically  very 
active.*  If  all  the  other  chemical  elements  differed  from  one  another 
as  much  as  do  these  three,  the  study  of  the  chemical  elements  would 
be  tedious  and  tiresome,  since  we  should  be  denied  the  satisfaction  of 
tracing  resemblances,  and  the  elements  would  be  incapable  of  classi- 
fication. In  reality,  however,  we  find  that  they  are  not  incapable  of 
being  grouped  together  in  sets.  They  are  classified  according  to  the 
kind  of  substances  with  which  they  combine  and  the  chemical  nature 
of  the  products.  In  some  families  the  resemblance  is  close,  in  others 

*  The  difference  between  oxides  and  chlorides  is  seen  in  their  behavior. 
Thus,  oxides  often  Unite  with  water  to  form  acids  or  bases  (p.  149).  Chlorides  do 
not  unite  with  water  to  form  new  substances  with  marked  characteristics  (cf.  p. 
152).  The  chlorides  of  metallic  elements  are  designated  salts  (p.  214). 

267 


268  INORGANIC  CHEMISTRY 

less  close.  The  present  group  is  of  the  former  class,  and  will  serve, 
therefore,  as  a  convenient  beginning  in  the  work  of  tracing  relations 
between  the  elements  and  in  classifying  the  facts  of  descriptive 
chemistry. 

The  Chemical  Relations  of  the  Halogens.  —  The  bromide 
(NaBr),  iodide  (Nal),  and,  to  a  less  extent,  the  fluoride  (NaF),  of 
sodium,  resemble  sodium  chloride  (NaCl)  in  composition,  appearance 
and  chemical  behavior.  From  this  fact,  chlorine,  bromine,  iodine, 
and  fluorine  are  known  as  the  halogens  (Gk.  o\s,  salt;  yivva.v,  to  pro- 
duce), and  their  compounds  are  named  the  halides.  The  halogens,  as 
the  above  formulae  show,  are  univalent.  They  all  form  compounds 
with  hydrogen,  and  these  compounds  closely  resemble  hydrogen  chlo- 
ride. For  example,  they  are  colorless,  they  are  gases  (with  the 
exception  of  hydrogen  fluoride  which  is  a  very  volatile  liquid),  they 
are  very  soluble  in  water,  and  their  solutions  are  acids.  Other  rela- 
tions will  be  given  in  a  summary  at  the  end  of  the  chapter. 

BROMINE 

Occurrence.  —  The  compounds  of  chlorine,  bromine,  and  iodine 
usually  occur  together  in  nature,  while  the  compounds  of  fluorine  are 
not  found  in  the  same  sources.  Bromine  occurs  chiefly  in  the  form  of 
the  bromides  of  sodium  and  magnesium,  in  the  upper  layers  of  the 
natural  beds  of  rock  salt.  Liebig  made  it  from  this  source  and  a  little 
later  Ballard  (1826)  made  it  also  and  recognized  it  as  a  new  element. 

Preparation.  —  In  the  chemical  point  of  view  there  are  three 
distinct  ways  in  which  bromine  is  made.  1 .  The  first  of  these  is  closely 
related  to  a  common  method  of  preparing  chlorine  (p.  220).  As 

hydrobromic  acid,  unlike  hydrochloric 
acid,  is  not  formed  extensively  in  connec- 
tion with  any  chemical  industry,  potas- 
sium bromide  is  treated  in  a  retort  (Fig. 

94)  with  concentrated  sulphuric  acid,  and 

FlG  94  the  product  is  oxidized  with  powdered 

manganese  dioxide  in  one  operation.    For 

the  equation,  see  next  section.  Bromine  being  a  volatile  liquid, 
while  the  sulphates  of  potassium  and  manganese  are  involatile, 
its  vapor  passes  off  when  the  above  mixture  is  heated.  It  is  con- 
densed in  a  flask  or  a  worm-tube  surrounded  by  cold  water. 

2.   The  second  method  of  preparing  bromine  depends  on  the  fact 


THE  HALOGEN  FAMILY  269 

that  chlorine  is  a  more  active  element  and  displaces  bromine  from 
combination.  When,  therefore,  chlorine  is  passed  into  a  solution  of 
potassium  or  sodium  bromide,  potassium  or  sodium  chloride  is  formed 
and  the  bromine  liberated: 

2NaBr  +  C12  ->  2NaCl  +  Br2. 

When  the  liquid  is  warmed,  the  bromine  passes  off  along  with  a  part 
of  the  water,  and  may  be  condensed  as  before. 

3.  Aqueous  solutions  of  soluble  bromides  may  be  decomposed  by 
means  of  a  current  of  electricity.  The  bromine  is  set  free  at  the 
positive  electrode. 

The  bromine  of  commerce  is  manufactured  in  the  first  two  of  these 
ways.  Two-thirds  of  the  world's  supply  is  obtained  from  Stassfurt, 
where,  after  the  extraction  of  the  potassium  chloride  from  the  impure 
carnallite  KCl,MgCl2,6H2O,  the  mother-liquor  is  found  to  contain 
the  more  soluble  sodium  and  magnesium  bromides  in  considerable 
quantities.  The  warm  mother-liquor  trickles  down  over  round  stones 
in  a  tower.  The  chlorine  is  introduced  from  below  and  dissolves  in 
the  liquid.  The  bromine  is  thus  liberated  and  passes  off  as  vapor. 
Our  supply  of  bromine  is  obtained  from  the  brines  of  Michigan, 
Ohio,  West  Virginia,  and  Kentucky.  Here  the  liquid,  after  most  of 
the  common  salt  has  been  removed  by  crystallization,  is  assayed  to 
ascertain  the  quantity  of  bromine  which  it  contains.  It  is  then 
treated  with  the  calculated  amount  of  sulphuric  acid  necessary  for  the 
action,  and  manganese  dioxide  is  added  gradually.  In  Michigan  the 
brines  are  treated  with  electrolytic  chlorine.  The  quantity  produced 
in  America  in  1914  was  288  short  tons. 

Partial  Equations.,  a  Plan  for  Making  Complex  Equations. 

—  When  an  equation  involves  more  than  two  initial  substances  or 
products,  as  does  the  one  given  above  for  the  first  method  of  preparing 
bromine,  it  cannot  readily  be  worked  out  by  the  method  formerly 
recommended  (p.  73).  After  the  formulae  of  all  the  substances,  on 
both  sides,  have  been  set  down,  it  is  difficult  to  hit  upon  the  proper 
numerical  factors  required  to  balance  the  equation.  In  such  cases 
a  good  plan  is  to  select  two  of  the  initial  substances  and  make  a  partial 
equation  showing  part  of  the  action  and  including  at  least  one  actual 
product.  Any  unused  units  (not  constituting  a  product)  are  then  set 
down  also  and  treated  as  a  balance.  Thus  the  first  two  of  the  sub- 
stances named  will  furnish  potassium-hydrogen  sulphate: 

Partial,  1 :  KBr  +  H2SO4  -» KHS04  (+  HBr).  (1) 


270  INORGANIC   CHEMISTRY 

Similarly,  the  manganese  dioxide  and  sulphuric  acid  will  give  man- 
ganous  sulphate: 

Partial,  2:       Mn02  +  H2S04  ->  MnS04  +  H20  (+  0).  (2) 

We  then  perceive  that  the  bromine  must  come  from  the  oxidation  of 
the  first  balance  by  the  second: 

Partial,  3 :  (2HBr)  +  (O)  -*  H20  +  Br2.  (3) 

The  third  partial  equation  shows  that  2HBr  will  be  needed  for  the 
amount  of  O  obtainable  from  MnO2,  so  we  go  back  to  (1)  and  multiply 
it  by  two  throughout: 

2KBr  +  2H2SO4  ->  2KHS04  (+  2HBr)  (1) 

MnO2  +  H2S04   -+  MnS04  +  H2O  (+  O)  (2) 

(2HBr)  +  (0)        ->  H20  +  Br2 (3) 

2KBr  +  3H2S04  +  Mn02   -»  2KHSO4  +  MnSO4  +  2H2O  +  Br2 

When  we  now  add  the  real  substances  used  and  produced,  as  they 
occur  in  these  partial  equations,  and  leave  out  the  balances,  which 
have  been  adjusted  so  as  to  cancel  one  another,  we  obtain  the  final 
equation  for  the  action.  It  must  be  observed  that  this  subdivision 
of  the  action  into  parts  is  a  purely  arithmetical  device,  used  solely  to 
simplify  the  arithmetical  process  of  writing  the  equations,  and  is  not 
intended  to  imply  that  the  chemical  change  itself  follows  these  or, 
indeed,  any  stages.  It  happens  that  the  three  partial  equations  we 
have  used  in  this  illustration  all  represent  interactions  which  can  take 
place  separately.  But  the  arithmetical  value  of  the  device  does  not 
depend  upon  this.  The  partial  equations  made  for  purposes  like  the 
present  one  are  often  purely  fictitious.  It  is  still  true,  however,  that 
we  are  aided  in  the  selection  of  partial  actions  at  each  step  by  follow- 
ing some  plausible  theory  as  to  stages  for  the  action  which,  if  there 
were  any,  would  be  chemically  conceivable. 

Physical  Properties.  —  Bromine  is  a  dark-red  liquid  (sp.  gr. 
3.18).  It  boils  at  59°,  forming  a  deep-red  vapor,  and  even  at  ordinary 
temperatures  gives  a  high  vapor  pressure  (150  mm.  at  18°)  and  evapo- 
rates quickly.  When  cooled  it  forms  red,  needle-shaped  crystals 
(m.-p.  —  7.3°).  A  saturated  aqueous  solution  (bromine-water)  at 
ordinary  temperatures  contains  3  parts  of  bromine  in  100  parts  of 
water.  The  element  is  much  more  soluble  in  carbon  bisulphide, 
alcohol,  and  other  organic  solvents.  Up  to  750°,  the  G.M.V.  weighs 
160  g.,  corresponding  to  Br2,  against  28.955  for  air. 


THE  HALOGEN  FAMILY  271 


Bromine  (Gk.  ftpupos,  a  stench)  has  a  most  pungent  odor.  It  has 
a  very  irritating  effect  on  the  mucous  membrane  of  the  nostrils  and 
throat.  If  spilled  upon  the  hands  it  has  a  most  destructive  action 
upon  the  tissues  and  leaves  sores  which  are  liable  to  infection. 

Free  bromine  has  no  effect  upon  starch  emulsion  (see  Iodine). 

Chemical  Properties.  —  The  molecules  of  bromine  are  less 
stable  than  are  those  of  hydrogen,  oxygen,  or  chlorine.  At  1050°  the 
G.M.V.  weighs  150.5,  and  dissociation  into  Br  has  begun. 

Bromine  unites  directly  with  hydrogen  to  form  hydrogen  bromide 
HBr,  but  the  mixture  of  the  gases  is  not  explosive,  and  the  union  is 
much  slower  than  in  the  case  of  chlorine. 

Bromine  forms  compounds  directly,  both  with  non-metals,  like 
phosphorus  and  arsenic,  and  with  most  of  the  metals,  which  catch 
fire  when  thrown  into  the  vapor.  Towards  unsaturated  substances 
and'organic  compounds  it  behaves  like  chlorine  (q.v.}.  In  all  cases 
the  interaction  is  less  violent  than  when  chlorine  is  used,  and  the 
element  is  displaced  from  combination  with  hydrogen  and  with  the 
metals  by  free  chlorine. 

Silver  bromide  is  the  sensitive  material  in  photographic  plates, 
and  potassium  and  sodium  bromides  are  used  as  sedatives  in  medicine. 
Bromine  is  required  in  large  quantities  in  the  manufacture  of  inter- 
mediate products  used  in  the  preparation  of  organic  dyes. 

HYDROGEN  BROMIDE  HBR 

Preparation.  —  It  might  be  expected  that  the  most  convenient 
way  of  producing  this  compound  would  be  similar  to  that  used  in  pre- 
paring hydrogen  chloride,  namely,  by  the  action  of  concentrated  sul- 
phuric acid  upon  some  common  bromide  such  as  potassium  bromide 
(KBr  +  H2SO4^  HBr  +  KHSO4).  We  find,  indeed,  that  at  first 
a  colorless  gas  is  given  off,  which  fumes  strongly  in  the  air  just  like 
hydrogen  chloride,  and  is  the  required  substance.  Almost  immedi- 
ately, however,  the  gas  acquires  a  yellow  and  then  a  brown  tinge,  and 
we  discover  that  free  bromine  is  being  produced  at  the  same  time.  If 
we  examine  the  gas  still  further,  we  recognize  also  the  presence  of 
sulphur  dioxide.  It  is  impossible,  therefore,  to  produce  hydrogen 
bromide  free  from  those  two  impurities  by  this  action. 

The  origin  of  the  bromine  and  sulphur  dioxide  which  complicate 
this  chemical  change  may  readily  be  traced.  Hydrogen  bromide  is 
less  stable  than  hydrogen  chloride,  and  its  hydrogen  Q&n  more  easily 


272 


INORGANIC  CHEMISTRY 


be  removed  by  the  action  of  substances  containing  oxygen.  In  this 
case  the  sulphuric  acid  acts  as  the  oxidizing  agent,  yielding  oxygen, 
sulphur  dioxide,  and  water  (H2SO4  ->  0  +  S02  +  H20).  Thus  the 
two  extra  gaseous  products  are  seen  to  be  formed  by  a  change  pro- 
ceeding parallel  with  the  main  action: 

2HBr  +  H2S04  -HJ  2H20  +  SO,T  +  Br2|  . 

The  simultaneous  occurrence,  in  this  fashion,  of  two  more  or  less 
independent  actions  in  the  same  vessel  is  not  uncommon.  Since  the 
HBr  is  first  liberated,  and  then  oxidized,  these  two  actions  are  called 
consecutive  actions.  The  speeds  of  such  actions  may  be  differently 
affected  by  temperature.  Thus,  here,  the  second  action  seems  to 
become  more  extensive  as  the  temperature  rises  (see  Chap.  XXIII). 

Since  all  acids  decompose  all  salts  more  or  less,  by  use  of  an  acid 
which  does  not  give  up  its  oxygen  so  readily,  such  as  phosphoric  acid, 
pure  hydrogen  bromide  may  be  obtained  (KBr  +  H3PO4  —  >  HBr  + 
KH2PO4)  .  The  small  solubility  of  the  salt  in  concentrated  phosphoric 
acid  retards  the  interaction  (p.  207)  and  makes  the  evolution  of  the 
gas  very  slow,  however. 

Pure  hydrogen  bromide  is  best  prepared  by  hydrolysis  (p.  210)  of 
phosphorus  tribromide.  When  bromine  and  phosphorus  are  mixed, 
a  violent  union  of  the  two  elements  takes  place,  producing  phosphorus 
tribromide  PBr3.  This  substance,  which  is  a  colorless  liquid,  is  in 
turn  broken  up  with  great  ease  by  water,  producing  phosphorous  acid, 
which  is  not  volatile,  and  gaseous  hydrogen  bromide: 

OH 


OH 

In  practice,  those  two  actions  are  carried  on  simultaneously.  To 

diminish  the  vigor  of  the  interaction,  red 
phosphorus  is  taken  instead  of  yellow,  and 
is  mixed  with  two  or  three  times  its  weight 
of  sand  in  a  flask  (Fig.  95).  A  small  quan- 
tity of  water  is  added.  Excess  of  water 
must  be  avoided,  as  the  hydrogen  bromide 
produced  is  extremely  soluble,  and  would 
therefore  be  retained  in  the  flask  instead 
of  being  disengaged  as  gas.  The  bromine 
is  placed  in  the  dropping  funnel,  and  ad- 

mitted, a  little  at  a  time,  to  the  mixture.    The  gas  produced  is 


FIG.  95. 


THE  HALOGEN  FAMILY  273 

passed  through  a  U-tube  containing  glass  beads  mixed  with  red  phos- 
phorus. The  latter  combines  with  any  bromine  which  may  have 
escaped  chemical  change  and  has  been  carried  along  with  the  gas. 
The  second  U-tube,  containing  water,  may  be  attached  when  a  solu- 
tion of  the  gas  is  required.  The  gas  may  be  collected  in  a  jar  by 
upward  displacement  of  air. 

Physical  Properties.  —  Hydrogen  bromide  is  a  colorless  gas 
with  a  sharp  odor.  It  is  two  and  a  half  times  as  heavy  as  air  (molar 
weight,  81  g.).  It  is  easily  reduced  to  the  liquid  condition  (b.-p. 
—  69°).  It  is  exceedingly  soluble  in  water,  and  in  contact  with  moist 
air  condenses  the  water  vapor  to  a  fog  of  liquid  particles.  When 
distilled,  the  solution  in  water  behaves  like  that  of  hydrogen  chloride 
(p.  211).  It  loses  mainly  either  water  or  hydrogen  bromide,  accord- 
ing as  it  is  dilute  or  exceedingly  concentrated,  until  an  acid  of  con- 
stant boiling-point  (126°  at  760  mm.  pressure),  containing  48  per  cent 
of  hydrogen  bromide,  passes  over.  Pure  hydrogen  bromide,  whether 
in  the  gaseous  condition  or  in  the  liquefied  form,  is  a  nonconductor  of 
electricity. 

Chemical  Properties.  —  The  chemical  properties  of  hydrogen 
bromide  are  similar  to  those  of  hydrogen  chloride  (p.  212) .  It  is  some- 
what less  stable,  and  dissociation  into  its  constituents  begins  to  be 
noticeable  at  800°.  When  free  from  water,  it  is  not  an  acid  (see 
below).  The  gas  interacts  vigorously  with  chlorine,  hydrogen 
chloride  and  free  bromine  being  produced,  and  much  heat  is  evolved 
by  the  change,  2HBr  +  C12  •**  2HC1  +  Br2.  The  heat  produced  by 
the  union  of  hydrogen  and  bromine  vapor  is  12,100  calories.  This  is 
much  less  than  the  amount  produced  by  the  union  of  chemically  equiv- 
alent quantities  of  hydrogen  and  chlorine  (22,000  calories).  When 
chlorine  displaces  bromine  from  hydrogen  bromide,  the  heat  evolved 
is  found  to  be  the  difference  between  these  two  numbers.  Using  the 
rule  of  constant  heat  summation  (p.  100),  we  write  equation  (2)  so  that 
HBr  is  on  the  same  side  with  Cl  (with  which  it  interacts),  and  the 
products  of  the  equation  required  (HC1  and  Br)  are  both  on  the  right: 

H  +  C1     ->HC1  +  22,000  cal.  (1) 

HBr          ->  Br  +  H     -  12,100  cal.  (2) 

Adding,  HBr  +  Cl  ->  HC1  +  Br  +    9900  cal. 

The  12,100  calories  are  produced  by  the  union  of  gaseous  bromine 
with  hydrogen,  and  the  final  result  is,  therefore,  that  for  the  produc- 


274  INORGANIC  CHEMISTRY 

tion  of  gaseous  bromine.  If  the  heat  of  formation  of  liquid  bromine 
is  required,  the  latent  heat  of  vaporization  of  bromine  (7296  calories), 
which  will  be  evolved  when  the  element  condenses,  must  be  added. 

Chemical  Properties  of  Hydrobromic  Acid  HBr,  Aq.  —  The 

solution  of  the  hydrogen  bromide  in  water  is  an  active  acid  (cf. 
p.  212).  It  conducts  electricity  extremely  well.  In  contact  with 
metals  above  hydrogen  in  the  order  of  activity  (p.  129),  and  with 
oxides  of  metals,  and  hydroxides  of  metals,  it  behaves  exactly  like 
hydrochloric  acid  (p.  213).  In  the  first  case,  hydrogen  is  set  free  and 
the  bromide  of  the  metal  produced.  In  the  other  two  cases,  water 
and  the  bromides  of  the  metals  are  produced.  For  example :  Zn(OH)2 
+  2HBr  — >  ZnBr2  +  2H20.  Oxidizing  agents  set  bromine  free  from 
hydrobromic  acid,  even  sulphuric  acid,  which  does  not  act  upon  hydro- 
chloric acid,  being  able  to  produce  this  result.  Chlorine  dissolved  in 
water  displaces  bromine  from  hydrobromic  acid  and  from  soluble 
bromides  with  ease  (test  for  bromides). 

IODINE  I2 

Occurrence.  —  Iodine,  like  bromine,  occurs  in  sea-water,  al- 
though in  much  smaller  quantities.  About  one-fifth  of  it  is  in  living 
algae  and  four-fifths  in  soluble  organic  compounds,  presumably  de- 
composition products  from  dead  algae,  and  little  as  mineral  iodides. 
Certain  species  of  sea-weed,  known  in  Scotland  as  kelp  and  in  Nor- 
mandy as  varec,  remove  it  from  water,  and  use  it  as  a  constituent  in 
complex  organic  compounds  which  they  contain.  The  ash  of  the 
sea-weed  sometimes  contains  as  much  as  two  per  cent,  or  even  more. 
The  chief  source  of  iodine,  however,  is  Chile  saltpeter  NaNO3,  in 
which  it  is  present  in  the  form  of  0.2  per  cent  of  sodium  iodate  NaI03. 
The  largest  proportion  of  iodine  in  the  human  body  is  in  the  thyroid 
gland.  In  diseases  like  goitre  and  cretinism,  where  the  thyroid  is 
ill-developed,  injection  of  a  substance  called  iodothyrine,  extracted 
from  sheep's  thyroids,  produces  marked  improvement. 

Preparation.  —  1.  In  factories  where  the  iodine  is  extracted 
from  sea-weed,  the  latter  is  carbonized  in  retorts.  The  residue  is 
extracted  with  water,  and  the  solution  is  evaporated  so  as  to  permit 
the  deposition  of  the  sodium  chloride  and  sodium  sulphate  which  it 
contains.  The  sodium  iodide,  being  very  soluble,  remains  in  the 
mother-liquor.  This  is  then  treated  with  manganese  dioxide  and 


THE  HALOGEN  FAMILY  275 

sulphuric  acid.  The  quantity  of  manganese  dioxide  is  carefully 
measured  so  as  to  be  just  sufficient  to  set  free  the  iodine  contained  in 
the  liquid,  without  proceeding  farther  to  the  liberation  of  the  chlorine 
which  it  contains  in  much  larger  amounts.  When  the  mixture  is 
heated,  the  iodine  passes  off  in  the  form  of  vapor,  and  is  condensed  in 
a  suitable  receiver.  The  action  (cf.  pp.  220,  270)  is: 

2NaI  +  Mn02  +  3H2S04  ->  MnS04  +  2NaHS04  +  2H20  +  I2  T  . 

2.  In  France  the  treatment  is  similar,  excepting  that  chlorine  is 
used  to  liberate  the  iodine  in  the  last  stage  (2NaI  +  C12  —  >  2NaCl  + 
I2).     The  quantity  is  adjusted  so  that  excess  may  not  be  employed. 
The  iodine,  being  insoluble,  forms  a  dense  precipitate  and,  when  the 
liquid  is  pressed  out,  it  remains  behind  in  the  form  of  a  paste. 

3.  Electricity  could  also  be  used  for  the  decomposition  of  this 
mother-liquor.     The  iodine  is  set  free  at  the  positive  electrode. 

4.  When  the  sodium  nitrate  has  been  crystallized  out  of  the 
aqueous  extract  from  the  Chile  saltpeter,  the  mother-liquor  is  treated 
with  sodium  sulphite  and  sodium  bisulphite: 

2NaI03  +  SNa^SOs  +  2NaHSO3  ->  5Na*S04  +  H20  +  I2  J,  . 

The  iodine,  being  insoluble,  is  precipitated. 

In  all  cases  the  iodine  is  purified  by  distillation  with  a  little 
powdered  potassium  iodide.  It  condenses  in  the  solid  form  directly, 
in  glittering,  black  plates  (sublimed  iodine).  The  distillation  of  a 
solid  body,  when  a  condensation  takes  place  directly  to  the  solid  form, 
is  spoken  of  as  sublimation. 


Physical  Properties.  —  Iodine  (Gk.  loei&js,  like  a  violet)  is  a 
solid  substance  (sp.  gr.  5),  exhibiting  large,  black  crystalline  plates  of 
rhombic  form.  It  melts  at  114°,  and  boils  at  184°.  The  vapor  has  at 
first  a  reddish-violet  tint  and,  on  being  more  strongly  heated,  be- 
comes deep  blue  (see  next  section). 

Iodine  is  very  slightly  soluble  in  water  (about  1:6000)  and  the 
solution  has  a  scarcely  perceptible  brown  tint.  It  is  much  more 
soluble  in  carbon  disulphide  (p.  189)  and  in  chloroform,  in  which  it 
gives  violet  solutions.  In  alcohol  and  ether  it  gives  solutions  which 
are  brown.  The  brown  color  is  attributed  to  the  fact  that  the  iodine 
is  in  a  condition  of  feeble  combination  and  not  simply  in  solution. 
These  liquids  become  violet  when  heated.  An  aqueous  solution  of 
potassium  iodide,  hydrogen  iodide,  or  any  other  iodide,  has  likewise 
the  power  to  take  up  large  quantities  of  iodine.  Here  the  formation 


276  INORGANIC  CHEMISTRY 

of  definite  compounds  (such  as,  KI  +  I2  <=±  KI3),  by  a  reversible 
action,  accounts  for  the  amount  of  iodine  which  appears  to  be  in 
solution. 

The  behavior  of  free  iodine  towards  starch  forms  a  distinctive  test 
for  both  substances  (cf.  p.  6).  When  the  pale-brown  aqueous  solu- 
tion, for  example,  is  added  to  a  filtered  starch  emulsion,  a  deep-blue 
color  is  produced.  The  same  action  is  used  as  a  test  for  starch.  This 
blue  material  is  not  a  chemical  compound.  The  iodine  is  adsorbed 
by  the  starch,  which  is  in  colloidal  suspension  (q.v.). 

Chemical  Properties.  —  The  molecular  weight  of  iodine,  as- 
certained by  weighing  the  vapor  at  temperatures  from  the  boiling- 
point  up  to  700°,  is  253.8.  The  atomic  weight  being  126.92,  the 
molecule  contains  two  atoms.  Beyond  700°,  the  vapor  diminishes  in 
density  more  rapidly  than  Charles'  law  would  lead  us  to  expect,  and 
at  1700°  the  molecular  weight  has  fallen  to  127  (cf.  p.  261).  As  the 
vapor  is  heated,  a  larger  and  larger  proportion  of  the  molecules  is 
broken  up,  until  the  decomposition  has  become  complete.  As  in  all 
cases  of  dissociation,  when  the  vapor  is  cooled  the  atoms  recombine 
to  form  molecules.  This  is  the  most  notable  case  in  which  we  en- 
counter both  the  monatomic  and  the  diatomic  forms  of  the  same 
element.  The  heat  given  out  when  the  atoms  reunite  to  form  the 
molecules  is  very  considerable  (21  +±  I2  +  28,500  cal.),  indicating  that 
the  chemical  union  of  two  atoms  of  identical  nature  may  be  as  vigor- 
ous as  that  of  two  atoms  of  different  chemical  substances.  The  heat 
of  union  of  atomic  hydrogen  (p.  253),  however,  is  even  greater 
(2H  <=^  H2  +  90,000  cal.  at  const,  press.).  The  monatomic  and 
diatomic  forms  of  iodine  should  be  distinct  chemical  substances,  and 
if  the  investigation  of  the  behavior  of  the  former  were  not  hampered 
by  the  very  high  temperature  at  which  alone  it  exists,  it  would  doubt- 
less be  found  to  exhibit  different  chemical  properties. 

Iodine  unites  very  slowly  with  hydrogen.  Iodine  unites  directly 
with  some  non-metals  and  with  the  majority  of  the  metals.  When 
phosphorus  is  presented  in  the  yellow  form,  the  action  takes  place 
spontaneously  without  the  assistance  of  heat.  Both  chlorine  and 
bromine  displace  iodine  from  combination  with  hydrogen  and  the 
metals  (2HI  +  Br2  — »  2HBr  +  I2) .  The  action  may  be  brought 
about  either  with  the  substances  in  dry  form  or  with  their  aqueous 
solutions. 

Iodine  in  water,  like  chlorine  in  water,  constitutes  an  oxidizing 
agent,  although  the  former  is  much  the  feebler  of  the  two.  The 


THE  HALOGEN  FAMILY  277 

hypo-iodous  acid,  formed  in  minute  amounts  at  a  time,  converts  sul- 
phurous acid  (q.v.)  into  sulphuric  acid: 

I2  +  H20     <=±HI  +  HIO. 
HIO  +  H2SO3  ->  HI  +  H2S04. 

In  analytical  chemistry  a  solution  of  iodine  in  potassium  iodide, 
containing  a  known  proportion  of  dissolved  iodine  (a  standard 
solution),  is  used  for  estimating  the  quantity  of  an  oxidizable  sub- 
stance present  in  a  given  specimen.  The  amount  of  oxidizable  sub- 
stance present  is  measured  by  the  quantity  of  the  standard  iodine 
solution  which  can  be  decolorized  and  suffer  removal  of  its  iodine. 
This  method  is  known  as  iodimetry. 

Iodine  and  its  compounds  are  much  used  in  the  arts  and  in  medi- 
cine. Iodine  is  applied  in  alcoholic  solution  (tincture  of  iodine)  for 
the  reduction  of  some  swellings,  and  as  an  antiseptic.  It  is  required 
in  making  iodoform  CHI3,  and  the  iodides  of  potassium,  rubidium, 
and  sodium,  which  are  used  in  medicine.  Hydriodic  acid  HI,  Aq  is 
used  to  promote  secretion  in  the  lymphatic  system.  The  emulsion 
used  in  making  photographic  dry-plates  contains  silver  iodide  Agl. 

HYDROGEN  IODIDE  HI 

Preparation.  —  The  direct  union  of  hydrogen  and  iodine  can- 
not be  employed  in  preparing  pure  hydrogen  iodide  (see  p.  306) . 
The  union  is  slow,  and  incomplete.  At  445°,  only  79  per  cent  of  the 
elements  unite,  because  the  action  is  reversible. 

The  action  of  concentrated  sulphuric  acid  upon  potassium  or 
sodium  iodide  is  equally  inapplicable.  In  this  case,  as  in  that  of 
hydrogen  bromide  (p.  272),  the  hydrogen  halide  reduces  the  sulphuric 
acid,  and  much  free  iodine  is  lormed.  Here,  on  account  of  the  greater 
ease  with  which  hydrogen  iodide  parts  with  its  hydrogen,  the  reduc- 
tion of  the  sulphuric  acid  is  much  more  complete,  the  product  being 
hydrogen  sulphide  H2S.  The  actions,  which  are  consecutive  (p.  272), 
are: 

KI  +  H2S04^HI  +  KHS04  and  H2S04  +  SHI"-*  H2S  +  4H20  +  4I2. 
As  soon  as  the  heat  produced  by  the  action  has  raised  the  temperature 
sufficiently,  hardly  any  of  the  hydrogen  iodide  escapes  oxidation.* 

*  When  much  sulphuric  acid  is  used,  sulphur  dioxide  and  free  sulphur  are 
formed  also.  This  is  in  consequence  of  a  secondary  action  of  the  hydrogen  sul- 
phide, as  it  passes  up  through  the  excess  of  sulphuric  acid,  and  of  the  sulphur 
dioxide  so  formed  upon  the  excess  of  hydrogen  sulphide  (q.v.): 

H2S  +  H2S04  -»  2H20  +  S  +  S02,          SO2  +  2H2S  ->  3S  +  2H2O. 


278  INORGANIC  CHEMISTRY 

Powdered  sodium  iodide  and  concentrated  phosphoric  acid  (cf. 
p.  272),  when  warmed,  give  pure  hydrogen  iodide  very  slowly: 
Nal  -f  H3PO4  <=±  HI  |  +  NaH2PO4.  This  action  was  formerly  used 
in  preparing  the  gas. 

The  best  method  is  one  similar  to  that  described  under  hydrogen 
bromide.  Phosphorus  and  iodine  unite  directly  to  form  PI3.  This  is 
a  yellow  solid  which  is,  violently  hydrolyzed  by  water  and  gives  phos- 
phorous acid  and  hydrogen  iodide: 

PI3  +  3H2O  -^  P(OH)3  +  SHI  T  . 

If  excess  of  water,  which  dissolves  hydrogen  iodide,  is  avoided,  the 
latter  goes  off  in  a  continuous  stream  in  a  gaseous  condition.  The 
mixture  of  iodine  (in  excess  of  the  amount  required  for  PI3)  and  red 
phosphorus  is  placed  in  the  flask  (Fig.  95)  and  the  water  in  the 
funnel. 

Still  another  method  of  making  hydrogen  iodide  is  frequently 
employed  when  a  solution  of  the  gas  in  water  is  required,  and  not  the 
gas  itself.  Powdered  iodine  is  suspended  in  water,  and  hydrogen  sul- 
phide gas  (q.v.)  is  introduced  through  a  tube  in  a  continuous  stream. 
The  iodine  dissolves  slowly  in  the  water,  I2  (solid)  «=^  I2  (dslvd),  and 
acts  upon  the  hydrogen  sulphide,  which  likewise  dissolves,  H2S  (gas)^ 
H2S  (dslvd)  .  Sulphur  separates  in  a  fine  powder,  S  (dslvd)  ^  S 
(solid),  and  hydrogen  iodide  is  formed  in  accordance  with  the  equa- 
tion: 

I2->2HI  +  S|  . 


This  action  takes  place,  however,  only  in  presence  of  water,  although 
the  water  does  not  appear  in  the  equation.  The  solution  is  freed  from 
the  deposit  of  sulphur  by  filtration,  and  may  be  concentrated  to  57 
per  cent  of  hydriodic  acid  by  distilling  off  the  water. 

The  theory  of  the  last  method  is  worthy  of  attention.  It  will  be 
seen  that  while  iodine  has  little  tendency  to  unite  with  free  hydrogen, 
it  is  here  able  to  decompose  a  compound  containing  hydrogen,  in  order 
to  secure  this  element.  It  is  enabled  to  do  this  by  the  fact  that  the 
very  large  amount  of  heat  given  out  by  the  mere  solution  of  hydrogen 
iodide  in  water  converts  the  action,  which  would  otherwise  be  endo- 
thermal,  into  an  exothermal  one.  In  the  absence  of  water,  the  reverse 
of  the  above  action  takes  place  with  ease.  In  presence  of  water, 
however,  the  great  heat  of  solution  (p.  202)  of  the  hydrogen  iodide 
(+19,200  cal.)  more  than  balances  the  heat  absorbed  by  the  chemical 


THE  HALOGEN  FAMILY  279 

change,  and  the  action  as  a  whole  takes  place  with  evolution  of  heat 
(see,  also,  Preparation  of  hydrogen  sulphide) : 

2HI  (dry)  +  S  ->  H2S  +  I2  +  14,200  cal. 
H2S,  Aq  + 12  ->  2HI,  Aq  +  S  +  19,600  cal. 

Physical  Properties.  —  Hydrogen  iodide  is  a  colorless  gas  with 
a  sharp  odor.  Its  molecular  weight  is  128,  and  it  is  therefore  much 
heavier  than  air,  the  average  weight  of  whose  molecules  is  28.955 
(p.  233).  It  is  a  nonconductor  of  electricity,  both  in  the  gaseous 
and  in  the  liquefied  conditions.  It  is  exceedingly  soluble  in  water,  so 
that  at  10°  thirty  grams  of  water  will  absorb  seventy  grams  of  the  gas, 
giving  a  70  per  cent  solution  (425  vols.  :  1  Aq).  The  behavior  of  this 
solution  is  similar  to  those  of  hydrogen  chloride  and  hydrogen  bro- 
mide (cf.  pp.  211,  273).  The  mixture  of  constant  boiling-point  distils 
over  at  127°  (at  760  mm.),  and  contains  57  per  cent  of  hydrogen  iodide. 

Chemical  Properties.  —  Hydrogen  iodide  is  the  least  stable  of 
the  hydrogen  halides.  When  heated  it  begins  visibly  (violet  color  of 
iodine  vapor)  to  decompose  into  its  constituents  at  180°.  On  account 
of  the  ease  with  which  it  parts  with  the  hydrogen  which  it  contains,  it 
can  bo  burned  in  oxygen  gas,  4HI  +  02  —>  2H2O  +  2I2.  When  the 
gas  is  mixed  with  chlorine,  a  violent  chemical  change,  accompanied  by 
a  flash  of  light,  occurs,  the  iodine  is  set  free,  and  hydrogen  chloride  is 
produced,  C12  +  2HI  — >  2HC1  +  I2.  Bromine  vapor  will  similarly 
displace  the  iodine  from  hydrogen  iodide. 

Chemical  Properties  of  Hydriodic  Acid  HI,  Aq.  —  In  most 
respects  the  aqueous  solution  behaves  exactly  like  hydrochloric  and 
hydrobromic  acids.  With  oxidizing  agents,  for  example,  such  as 
manganese  dioxide,  it  gives  free  iodine,  just  as  the  others  give  free 
chlorine  and  bromine,  respectively.  Here,  however,  the  oxidation  is 
so  much  more  easily  carried  out,  that  it  is  slowly  effected  by  atmos- 
pheric oxygen,  so  that  hydriodic  acid  left  exposed  to  the  air  gradually 
becomes  brown:  O2  +  4HI  — >  2H2O  +  2I2.  The  free  iodine  remains 
dissolved  in  the  hydrogen  iodide,  in  the  form  of  the  compound  HI3. 
Finally,  however,  the  free  iodine,  as  its  quantity  becomes  greater,  and 
that  of  the  hydrogen  iodide  smaller,  is  deposited  in  crystalline  con- 
dition. On  account  of  the  ease  with  which  hydriodic  acid  parts  with 
its  hydrogen,  it  is  frequently  used  in  chemistry  as  a  reducing  agent. 

Although  the  dry  gas  is  not  an  acid,  the  solution  has  all  the  ordi- 
nary properties  of  this  class  of  substances  (cf.  p.  117).  The  hydrogen 


280 


INORGANIC  CHEMISTRY 


may  be  displaced  by  metals  like  zinc  and  magnesium  (p.  118).  The 
acid  interacts  with  oxides  and  hydroxides,  forming  iodides  and  water 
(p.  213). 


FLUORINE  F2 

The  discussion  of  this  element  should  logically  have  preceded  that 
of  chlorine,  since  it  is  of  all  the  members  of  the  halogen  family  the 
most  active.  Chlorine  was  taken  up  first,  however,  because  its  com- 
pounds are  more  familiar.  Fluorine  is  found  in  nature  chiefly  in  the 
mineral  fluorite  (calcium  fluoride)  CaF2,  in  cryolite,  a  double  fluoride 
of  aluminium  and  sodium  AlF3,3NaF,  and  in  apatite  (q.v.).  It  is 
found  also  in  small  amounts  in  bones  (especially  the  teeth). 

Preparation.  —  When  a  solution  of  hydrofluoric  acid  is  heated 
with  manganese  dioxide,  oxidation  does  not  occur  and  free  fluorine 
is  not  produced.  Until  1886  all  efforts  to  isolate  the  element 

failed.  It  was  perfectly  understood  that  the 
reason  of  these  failures  lay  in  the  greater 
chemical  activity  of  fluorine,  which  made  it 
more  difficult  of  separation  from  any  state 
of  combination  than  the  other  halogens. 
Its  preparation  was  finally  achieved  by 
Moissan  (1886)  by  the  decomposition  of 
3  anhydrous  hydrogen  fluoride,  which  is  liquid 
below  19°,  by  means  of  electricity.  The 
apparatus  (Fig.  96)  is  made  of  copper,  which, 
after  receiving  a  thin  coating  of  the  fluoride, 
is  not  further  affected.  To  reduce  the 
tendency  to  chemical  union,  the  whole  is 
immersed  in  a  bath  giving  a  temperature 
of  —23°.  The  electrodes  are  made  of  an 
alloy  of  platinum  and  iridium,  which  is  the 
only  substance  that  can  resist  the  action  of 
the  fluorine  when  freshly  liberated  by  the 
electric  current.  Hydrogen  fluoride,  like 
other  hydrogen  halides,  is  a  nonconductor  of  electricity,  and  a  small 
quantity  of  potassium-hydrogen  fluoride  KHF2  has  to  be  added  to  en- 
able the  current  of  electricity  to  pass.  The  fluorine  is  set  free  at 
the  positive  electrode,  and  hydrogen  appears  at  the  negative.  The 
U-*ube  is  closed  after  the  introduction  of  the  hydrogen  fluoride  by 


FIG.  96. 


THE  HALOGEN  FAMILY  281 

means  of  blocks  made  of  calcium  fluoride,  which  is  naturally  unable 
further  to  enter  into  combination  with  fluorine.  For  the  reception 
and  examination  of  the  fluorine  gas,  other  copper  tubes  can  be 
screwed  on  to  the  side  necks  of  the  apparatus,  and,  when  necessary, 
small  windows  of  calcium  fluoride  can  be  provided.  It  has  been 
found  that  fluorine  dried  with  extraordinary  precautions  is  without 
action  on  clean,  dry  glass. 

Physical  Properties.  —  Fluorine  is  a  gas  whose  color  is  like 
that  of  chlorine,  but  somewhat  paler.  Its  density  indicates  a  molec- 
ular weight  of  38,  showing  that  the  molecule  is  diatomic  (the  atomic 
weight  is  19).  The  gas  is  the  most  difficult  of  the  halogens  to  liquefy 
(b.-p.  -186°). 

Chemical  Properties.  —  Fluorine  unites  with  every  element, 
with  the  exceptions  of  oxygen,  chlorine,  nitrogen,  and  the  helium 
family,  and  in  many  cases  does  so  with  such  vigor  that  the  union 
begins  spontaneously  without  the  assistance  of  external  heat.  Dry 
platinum  and  gold  are  the  elements  least  affected.  It  explodes  with 
hydrogen  at  the  ordinary  temperature,  without  the  assistance  of 
sunlight.  On  the  introduction  of  a  drop  of  water  into  a  tube  of 
fluorine,  the  oxygen  of  the  water  vapor  is  instantly  displaced  by  the 
fluorine,  and  the  vessel  is  filled  with  the  deep-blue  gas,  ozone:  3F2  + 
3H2O-+3H2F2  +  O3. 

The  chlorine  in  hydrogen  chloride  is  displaced  by  fluorine  as  easily 
as  chlorine  in  turn  displaces  bromine  or  iodine. 

HYDROGEN  FLUORIDE  H2F2 

Preparation.  —  Pure,  dry  hydrogen  fluoride  is  best  made  by 
heating  potassium-hydrogen  fluoride,  2KHF2  +±  2KF  +  H2F2  T . 
For  ordinary  purposes,  however,  the  preparation  of  an  aqueous  solu- 
tion is  the  ultimate  object.  Usually  powdered  calcium  fluoride  is 
treated  with  concentrated  sulphuric  acid,  and  the  mixture  distilled  in 
a  retort  of  platinum  or  lead: 

CaF2  +  H2S04  ?±  CaS04  +  H2F2  T . 

The  hydrofluoric  acid  passes  over  and  is  caught  in  distilled  water. 
The  aqueous  solution  thus  obtained  has  to  be  kept  in  vessels  made  of 
lead,  rubber,  or  paraffin,  as  glass  interacts  with  the  acid  with  great 
rapidity  (see  below). 


282  INORGANIC  CHEMISTRY 

Physical  Properties.  —  Hydrogen  fluoride  is  a  colorless  liquid, 
boiling  at  19.4°.  It  mixes  freely  with  water  and,  on  distillation,  an 
acid  of  constant  boiling-point  (120°  at  760  mm.)  containing  35  per 
cent  of  hydrogen  fluoride  is  obtained.  The  weight  of  22.4  liters  of 
the  vapor  varies  from  51  g.  at  26°  to  20  g.  at  90°  and  above.  At  90° 
and  above,  therefore,  the  formula  is  HF,  and  at  26°  the  vapor  is 
mainly  a  mixture  of  H2F2  (40)  and  H3F3  (60).  Since  HF  is  the  only 
form  which  persists  through  a  range  of  temperatures,  we  say  that  the 
substance  shows  association  (see  below)  at  lower  temperatures.  To 
keep  ourselves  in  mind  of  this  peculiarity,  we  shall  use  the  formula 
H2F2,  although  the  liquid  undoubtedly  contains  many  molecules 
which  are  higher  multiples  of  HF  than  this. 

Association.  —  Many  substances  resemble  hydrogen  fluoride  in 
consisting  of  mixtures  of  molecules  which  are  multiples  of  the  simplest 
possible.  Thus,  acetic  acid  vapor  at  the  boiling-point  has  the  for- 
mula (HCO2CH3)2,  and  sulphur  vapor  at  high  temperatures  is  S2  (p. 
261),  but  at  lower  temperatures  a  mixture  of  82,  Se,  and  Sg.  These  are 
associated  vapors. 

There  are  also  associated  liquids.  Thus,  sulphuric  acid  and  nitric 
acid  in  the  liquid  condition  are  composed  of  more  complex  aggregates 
than  H2SO4  and  IINO3.  Even  water  is  largely  (H2O)2  or  even  (H2O)3, 
although  the  vapor  is  H2O  (p.  202) .  In  all  such  cases  dissociation  into 
the  simpler  molecules  takes  place  gradually  as  the  temperature  is 
raised. 

Many  substances  naturally  possess  formulae  which  are  multiples  of 
the  simplest  without  showing,  as  the  temperature  is  raised,  any  tend- 
ency to  progressive  dissociation  into  the  corresponding  simplest 
molecules.  Thus,  acetylene  (p.  241)  is  C2H2  at  all  temperatures,  and 
acetic  acid  (q.v.),  although  it  is  associated  to  C^aCh  at  its  boiling- 
point,  never  becomes  simpler  than  C2H4O2  at  any  temperature. 

When  a  substance  changes  sharply  into  another  substance  with 
double  or  triple  molecular  weight,  as  formaldehyde  CH20,  a  volatile 
liquid,  changes  into  para-formaldehyde  (CH2O)3,  a  crystalline  solid, 
the  phenomenon  is  called  polymerization. 

Chemical  Properties  of  Hydrofluoric  Acid.  —  Metals  like  zinc 
and  magnesium  interact  with  hydrofluoric  acid  with  evolution  of 
hydrogen  (p.  118).  The  action  is  less  violent  than  with  other  halogen 
acids.  The  acid  interacts  with  oxides  arid  hydroxides,  forming 
fluorides  (p.  213).  The  chief  difference  in  this  respect  which  it  ex- 


THE  HALOGEN  FAMILY  283 

hibits,  when  compared  with  the  other  halogen  acids,  is  the  one  which 
leads  us  to  assign  to  it  the  formula  H2F2.  We  may  displace  either 
one  or  both  the  hydrogen  atoms  in  the  molecule  with  a  metal.  Thus, 
one  of  the  commonest  salts  of  hydrofluoric  acid  is  potassium-hydrogen 
fluoride  HKF2,  mentioned  above,  KOH  +  H2F2  ->  KHF2  +  H2O. 
In  this  respect  the  acid  resembles  sulphuric  acid  and  other  acids 
containing  more  than  one  replaceable  hydrogen  unit.  Unlike  the 
other  halogen  acids,  hydrofluoric  acid  gives  acid  salts  with  great  ease. 
The  most  remarkable  property  of  hydrofluoric  acid  is  that  it 
interacts  readily  with  sand,  silicon  dioxide  SiO2,  giving  silicon  tetra- 
fluoride  SiF4,  a  gas,  and  water: 

Si02  +  2H2F2  -p>  SiF4  T  +  2H2O. 

No  other  acid  is  able  thus  to  act  upon  the  oxide  of  a  typical  non- 
metallic  element.  In  fact,  in  all  other  cases  (e.g.,  SiCU),  water 
decomposes  the  halide  (hydrolysis),  and  the  action  goes  in  the  opposite 
direction. 

Glass,  which  is  commonly  made  by  fusing  together  sodium  car- 
bonate, calcium  carbonate,  and  sand  (silicon  dioxide),  is  a  mixture  of 
silicates  of  calcium  and  sodium,  and  is  rapidly  decomposed  by  hydro- 
fluoric acid.  The  nature  of  the  change  is  shown  by  the  two  following 
equations  : 

CaSi03  +  3H2F2-*SiF4  1  +  CaF2  +  3H2O. 
3H2F2  ->  SiF4  f  +  2NaF  +  3H20. 


All  other  silicates  are  decomposed  according  to  the  same  plan.  The 
silicon  tetrafluoride  passes  off.  The  fluorides  of  calcium  and  sodium 
are  solid  and  crumble  away  or  dissolve.  Thus  the  glass  is  completely 
disintegrated.  The  vapor  of  hydrofluoric  acid,  generated  in  the  way 
described  above  from  calcium  fluoride  in  a  lead  dish,  is  used  for  etch- 
ing glass.  The  surface  of  the  glass  is  covered  with  paraffin  to  protect 
it  from  the  action  of  the  vapor,  and  with  a  sharp  instrument  portions 
of  this  paraffin  are  removed  where  the  etching  effect  is  desired.  The 
vapor  gives  a  rough  surface  where  it  encounters  the  glass  (test  for  a 
fluoride)  .  Burettes  and  other  glass-ware  are  graduated  in  this  fashion. 
The  aqueous  solution,  which  may  also  be  employed,  makes  smooth 
depressions  on  the  surface.  The  aqueous  solution  of  the  acid  is  used 
in  the  analysis  of  minerals  containing  silicates,  which  frequently  are 
not  attacked  by  other  acids.  It  is  used  also  for  removing  sand  from 
metal  castings  and  for  cleaning  the  exteriors  of  buildings  of  granite 
and  sandstone. 


284 


INORGANIC  CHEMISTRY 


THE  HALOGENS  AS  A  FAMILY 

The  most  noticeable  fact  is  that,  if  we  arrange  the  halogens  in 
order  in  respect  to  any  one  property,  chemical  or  physical,  the  other 
properties  will  be  found  to  place  them  in  the  same  order.  In  the 
table,  below,  the  sixth  column  contains  the  weight  of  the  element 
in  liquid  or  solid  form  dissolving  in  100  c.c.  of  water  (15°).  The 
last  column,  cal.  KX,  gives  the  heat  of  formation  of  one  gram- 
molecule  of  the  potassium  halide. 


Element. 

At.  Wt. 

State. 

B.-p. 

Color. 

Sol'ty. 

Cal.  KX 

Fluorine      
Chlorine  

19.0 
35.5 

gas 
gas 

—187° 
—  34 

yellow 
yellow 

118,100 
104,300 

Bromine      

79.9 

liq. 

+    59 

brown 

3.22 

95,100 

Iodine      

126.9 

solid 

+184° 

violet 

0.015 

80,100 

It  will  be  seen  that,  as  the  atomic  weight  (chem.  prop.)  increases,  the 
boiling-point  (b.-p.)  rises,  the  color  deepens  and  passes  towards  the 
blue  end  of  the  spectrum,  the  solubility  diminishes  (phys.  props.), 
and  the  heat  of  union  with  potassium  (chem.  prop.)  becomes  smaller. 
The  vigor  with  which  the  halogens  unite  with  hydrogen  and  the  metals 
is  greatest  with  fluorine  and  diminishes  progressively  until  we  reach 
iodine.  We  shall  see  later  that  the  affinity  for  oxygen,  on  the  other 
hand,  increases  as  we  pass  from  fluorine  to  iodine. 

Although  showing  different  degrees  of  activity,  the  halogens  are 
closely  alike  in  chemical  nature.  That  is,  the  relations  (p.  226)  they 
show  when  in  combination  are  similar.  When  united  with  hydrogen 
and  the  metals,  they  are  all  univalent.  In  their  oxygen  compounds, 
however,  they  exhibit  a  higher  valence.  Their  oxides  interact  with 
water  to  give  acids,  and  they  are  therefore  non-metals  (p.  150).  They 
are  strongly  electro-negative  (pp.  213,  269),  as  non-metals  all  are. 
Their  hydrides,  when  dissolved  in  water,  are  all  active  acids.  This, 
and  their  valence,  distinguish  the  halogen  family  from  other  groups 
of  non-metals.  Thus,  oxygen  and  sulphur  are  bivalent  (and  the  latter 
sexivalent  also),  and  the  hydrides  of  oxygen  (H2O  and  H2O2)  and  of 
sulphur  (£[28)  are  very  feeble  acids. 

Order  of  Activity  of  the  Non-Metals.  —  The  way  in  which 
chlorine  displaces  bromine  and  iodine  from  bromides  (p.  273)  and 
iodides  (p.  276),  and  bromine,  in  turn,  displaces  iodine  suggests  an 


THE  HALOGEN  FAMILY  285 

order  of  activity  for  non-metals.  It  was  noted  that  oxygen  displaced 
iodine  from  bydriodic  acid  (p.  279)  and  that  iodine  displaced  sulphur 
from  hydrogen  sulphide  (and  all  other  sulphides).  The  order  is, 
therefore,  F,  Cl,  Br,  0,  I,  S. 

COMPOUNDS  OF  THE  HALOGENS  WITH  EACH  OTHER 

Iodine  unites  with  chlorine  to  form  two  compounds.  The  most 
familiar  is  a  red  crystalline  substance  iodine  monochloride  Id.  The 
trichloride  IC13  is  made  by  the  use  of  excess  of  chlorine.  Iodine 
unites  with  bromine  to  form  the  compound  IBr,  while  a  compound 
with  fluorine  IFs,  is  supposed  to  exist.  None  of  these  compounds  is 
particularly  stable,  and  some  of  them  decompose  very  easily. 

It  is  frequently  stated  that  elements  which  resemble  one  another 
chemically  show  little  tendency  to  chemical  union.  Yet  in  the  case 
of  IBr,  for  example,  the  tendency  to  decompose  into  the  elements 
(2IBr  — >  I2  +  Br2)  must  be  interpreted  as  meaning  that  the  iodine 
and  bromine  prefer  to  unite  with  themselves  to  form  the  molecules 
I2  and  Br2  rather  than  with  one  another.  In  view  of  this  the  above 
remark  loses  some  of  its  point,  for  an  element  certainly  resembles  itself 
more  than  it  does  any  other,  and  the  compounds  C12,  H2,  etc.,  are 
amongst  the  most  stable  that  we  know. 

Exercises.  —  1.  What  impurities  is  commercial  iodine  likely  to 
contain?  In  what  way  does  heating  with  potassium  iodide  (p.  275) 
free  it  from  these? 

2.  Classify  all  the  chemical  actions  in  this  chapter  according  as 
they  belong  to  one  or  other  of  the  ten  kinds  (p.  228). 

3.  What  are  the  relative  volumes  of  the  gases  in  the  interaction  of 
chlorine  with  hydrogen  bromide  (p.  273),  and  hydrogen  iodide  (p. 
279),  respectively? 

4.  Tabulate,  more  fully  and  specifically  than  is  done  in  the  section 
on  "The  Halogens  as  a  Family,"  (a)  the  physical  properties,  (b)  the 
chemical  properties,  (c)  the  chemical  relations,  of  the  members  of 
this  group. 

5.  Construct  the  equation  on  p.  275  by  the  use  of  partial  equations 
as  in  the  example  on  p.  270. 

6.  Using  the  method  given  on  p.  270,  construct  a  single  equation 
for  the  formation  of  iodine,  water,  and  hydrogen  sulphide  directly 
from  potassium  iodide  and  sulphuric  acid. 

7.  What  are  the  relative  volumes  of  the  gases  in  the  action  of, 


286  INORGANIC  CHEMISTRY 

(a)  fluorine  and  water  vapor  (p.  281),  (6)  chlorine  and  iodine  vapors 
in  forming  the  monochloride  and  (c)  the  trichloride. 

8.  In  the  French  process  for  liberating  iodine  (p.  275),  why  must 
excess  of  chlorine  be  avoided? 

9.  In  liberating  bromine  from  the  mother-liquor  (p.  269),  why 
must  excess  of  sulphuric  acid  and  manganese  dioxide  be  avoided? 


CHAPTER  XV 
CHEMICAL  EQUILIBRIUM 

IN  spite  of  its  formidable  title,  this  chapter  will  introduce  nothing 
novel.  Its  purpose  is  to  collect  together  and  organize  more  definitely 
a  number  of  scattered  facts  and  ideas  which  have  already  come  up  in 
various  connections.  On  this  account,  however,  it  will  be  all  the  more 
necessary  for  the  reader  to  refresh  his  remembrance  of  these  facts 
and  ideas  by  re-reading  all  pages  to  which  reference  is  made. 

Reversible  Actions.  —  In  discussing  the  union  of  hydrogen  and 
iodine  (p.  277),  it  was  stated  that  the  progress  of  the  action  ceases 
while  yet  a  large  amount  (21  per  cent  at  445°)  of  both  the  substances 
necessary  for  its  maintenance  still  remains  available.  Now  the 
materials  left  over  are  presumably  no  less  capable  of  uniting  than  the 
parts  which  have  already  united.  The  solution  of  this  mystery  lies 
in  the  fact  (p.  279)  that  decomposition  of  the  compound  can  begin  at 
180°,  and  therefore  takes  place  actively  at  445°.  Hence  the  product 
of  the  union  must  begin  to  dissociate,  in  part  at  least,  as  soon  as  any 
of  it  is  formed.  Thus  two  changes,  one  of  which  undoes  the  work  of 
the  other,  must  go  on  simultaneously.  In  consequence  of  this,  neither 
can  reach  completion.  As  we  should  expect,  experiment  shows  that 
it  makes  no  difference  whether  we  start  with  pure  hydrogen  iodide  or 
with  a  mixture  of  pure  hydrogen  and  iodine:  the  proportions  of  the 
three  materials  found  in  the  tube,  after  it  has  been  heated  for  a 
sufficient  length  of  time,  are  in  both  cases  the  same.  A  general 
statement  may  be  founded  on  facts  like  this,  to  the  effect  that  a 
chemical  action  must  remain  more  or  less  incomplete  when  the  re- 
verse action  also  takes  place  under  the  same  conditions  (cf.  p.  40). 
Two  arrows  pointing  in  opposite  directions  are  used  in  equations 
representing  reversible  changes: 

H2  +  I2<=±2HI    or    2HI<=»H2  +  I2. 

It  will  be  observed  that  representing  reversible  actions  by  equations  involves 
a  departure  from  the  original  meaning  of  an  equation.  Thus  at  445°,  79  per  cent 
of  the  substances  are  in  the  form  HI  and  21  per  cent  in  the  uncombined  state: 

(79%)2HI<=±H2+I2(21%). 
287 


288  INORGANIC  CHEMISTRY 

In  other  words,  the  amounts  of  matter  on  the  two  sides  are  not  equal.  Each  side, 
taken  separately,  shows  correctly  the  proportions  used  in  the  interaction  for  which 
it  stands,  however.  Hence  the  equation  in  a  reversible  action  professes  to  show 
quantitatively  the  change  which  would  occur  if  each  of  the  two  opposed  actions  it 
includes  were  to  be  allowed  separately  to  proceed  to  completion. 

The  following  are  examples  of  actions  of  the  exactly  same  kind. 
They  should  now  be  looked  up  and  studied  attentively.  The  dis- 
cussion in  the  following  sections,  for  which  they  furnish  the  basis, 
cannot  otherwise  be  understood:  (1)  The  interaction  of  chlorine  and 
water  (p.  223),  which  was  fully  discussed  at  the  time;  (2)  the  be- 
havior of  iodine  vapor  (p.  261),  of  water  vapor  (p.  148),  of  sulphur 
vapor  (p.  261),  of  phosphorus  vapor  (p.  261);  (3)  the  behavior  of 
phosphorus  pentachloride  vapor  (p.  260) ;  (4)  Deacon's  reaction  (p. 
217).  Examples  of  this  kind  are  called  homogeneous  systems,  be- 
cause all  the  interacting  substances  are  either  gaseous  or  in  solution 
(C12  +  H20). 

When  the  action  is  one  which  is  reversible,  but,  under  the  circum- 
stances being  discussed,  proceeds  farther  towards  completion  in  one 
direction  than  in  the  other,  the  arrow  will  be  modified  to  indicate  this 
fact: 

C12  +  H2O  <=»  HC1  +  HC1O  (p.  223). 

When  this  relative  completeness  is  due  to  volatilization  or  pre- 
cipitation, the  fact  will  be  indicated  by  vertical  arrows: 

NaCl  +  H2S04^NaHS04  +  HC1  T  (p.  207). 

NaCl  |  +  H2S04  *=?  NaHS04,  Aq  +  HC1,  Aq  (p.  208). 

Cases  in  which  a  gas  or  a  solid  separates  are  called  heterogeneous 
systems. 

Actions  which  Proceed  to  Completion.  —  All  chemical  ac- 
tions do  not  belong  to  the  reversible,  incomplete  class.  Many  pro- 
ceed uninterruptedly  to  exhaustion  of  one,  or  all,  of  the  ingredients. 
For  example,  equivalent  amounts  of  magnesium  and  oxygen  combine 
completely:  2Mg  +  O2  — >  2MgO.  Here,  however,  the  product  is  not 
decomposed  even  at  the  white  heat  produced  by  the  vigor  of  the  union. 
Indeed,  magnesium  oxide  cannot  be  decomposed,  and  the  action  re- 
versed, at  any  temperature  we  can  command.  The  other  complete 
actions,  like  the  decomposition  of  potassium  chlorate  (p.  83),  are  so 
because  they  are  likewise  irreversible. 


CHEMICAL  EQUILIBRIUM  289 

Kinetic- Molecular  Explanation.  —  Restating  these  facts  in 
terms  of  molecules  will  enable  us  to  reason  more  clearly  about  this 
variety  of  chemical  change.  Suppose  we  start  with  the  materials 
represented  on  one  side  only  of  such  an  equation  say  the  hydrogen 
and  iodine  in  that  on  p.  287.  The  molecules  of  these  materials  will 
encounter  one  another  frequently  in  the  course  of  their  movements. 
In  a  certain  proportion  of  these  collisions  the  chemical  change  will 
take  place.  In  the  earliest  stages  there  will  be  few  of  the  new  kind  of 
molecules  (say,  hydrogen  iodide),  but,  as  the  action  goes  on,  these  will 
increase  in  number.  There  will  be  two  consequences  of  this.  In  the 
first  place  the  parent  materials  (in  this  case,  hydrogen  and  iodine)  will 
diminish  in  amount,  the  collisions  between  their  molecules  will  be- 
come fewer,  and  the  speed  of  the  forward  action  will  therefore  become 
less  and  less.  In  the  second  place  the  increase  in  the  number  of 
molecules  of  the  product  (hydrogen  iodide)  will  result  in  more  frequent 
collisions  between  them,  in  more  frequent  occurrence  of  the  chemical 
change  which  they  can  undergo,  and  thus  in  an  increase  in  the  speed 
of  the  reverse  action.  The  forward  action  begins  at  its  maximum  and 
decreases  in  speed  progressively;  the  reverse  action  begins  at  zero  and 
increases  in  speed.  Finally  the  two  speeds  must  become  equal,  and 
at  that  point  perceptible  change  in  the  condition  of  the  whole  must 
cease  (cf.  pp.  169-170). 

The  most  immediate  inference  from  this  mode  of  viewing  the 
matter  is,  that  the  apparent  halt  in  the  progress  of  the  action  does 
not  indicate  any  cessation  of  either  chemical  change.  Both  changes 
must  go  on,  in  consequence  of  the  continued  encounters  of  the  proper 
molecules.  But  since  the  two  changes  proceed  with  equal  speeds  they 
produce  no  alteration  in  the  mass  as  a  whole.  In  fact,  the  final  state 
is  one  of  equilibrium,  and  not  of  repose.  Hence,  chemical  changes 
which  are  reversible  lead  to  that  condition  of  seemingly  suspended 
action  which  we  speak  of  as  chemical  equilibrium. 

Chemical  Equilibrium  and  its  Characteristics.  —  The  de- 
tailed discussions  of  the  relations  of  liquid  and  vapor  (pp.  145,  169- 
170),  of  ice  and  water  (p.  144),  of  saturated  solution  and  undissolved 
solid  (pp.  185,  192-195),  and  of  a  gas  dissolving  in  a  liquid  (p.  187), 
have  already  familiarized  us  with  the  term  equilibrium  and  its 
significance.  We  can,  in  fact,  apply  to  the  discussion  of  any  kind 
of  reversible  phenomena,  the  sets  of  ideas  in  regard  to  exchanges  of 
molecules  there  elaborated. 

In  particular,  the  reader  will  note  that  the  three  characteristics  of 


290  INORGANIC   CHEMISTRY 

a  state  of  equilibrium,  developed  and  illustrated  in  the  case  of  the 
physical  equilibrium  between  a  liquid  and  its  vapor  (p.  169),  apply 
also  to  a  typical  case  of  chemical  equilibrium,  such  as  that  now  before 
us.  Thus : 

1.  There  are  the  two  opposing  tendencies,  which  ultimately  bal- 
ance one  another.     Here  they  are  the  tendency  of  the  hydrogen  and 
iodine  to  produce  hydrogen  iodide,  and  the  tendency  of  the  hydro- 
gen iodide  to  reproduce  hydrogen  and  iodine  by  this  interaction.     In 
other  words,  they  are  the  apparent  activity  of  the  hydrogen  iodide 
reaction,  and  the  apparent  activity  *  of  the  hydrogen  and  iodine 
interaction. 

2.  At  equilibrium  the  two  opposing  tendencies  or  activities  are  still 
in  full  operation,  although  their  effects  then  neutralize  one  another. 

3  (and  this  is  the  chief  mark  of  chemical,  as  it  is  of  physical 
equilibrium).  The  system  is  in  a  sensitive  state,  so  that  a  change 
in  the  conditions  (temperature  and  partial  pressure  or  concentration), 
even  if  slight,  produces  a  corresponding  change  in  the  state  of  the 
system,  and  does  this  by  favoring  or  disfavoring  one  of  the  two  oppos- 
ing tendencies  or  apparent  activities.  Such  a  change  is  called  a  dis- 
placement of  the  equilibrium,  for  the  system  settles  down  in  a  new 
state  of  equilibrium  with  new  proportions  of  the  two  sets  of  sub- 
stances, corresponding  to  the  changed  conditions.  Thus,  in  the 
present  instance,  a  change  from  445°,  where  there  is  79  per  cent  of 
the  material  in  the  form  of  hydrogen  iodide,  to  508°  results  in  the 
diminution  of  this  proportion  to  76  per  cent.  The  equilibrium  is 
affected  by  changes  in  concentration  also,  as  we  shall  presently  see 
(pp.  291-293). 

Now,  the  foregoing  facts  show  that  the  key  to  understanding 
chemical  activities,  their  magnitudes,  their  changes,  and  especially 
their  practical  results,  must  lie  in  knowing  how  changes  in  the  condi- 
tions affect  them.  Hence,  to  the  chemist,  familiarity  with  the  in- 
fluence of  conditions  on  chemical  phenomena  must  be  of  the  greatest 
practical  importance.  We  therefore  address  ourselves  to  the  dis- 
cussion of  this  subject. 

The  "conditions"  to  be  considered  are  familiar  —  temperature, 
and  concentration  or,  in  the  case  of  a  gas,  partial  pressure.  The 
"activity"  of  an  action  is  accurately  measured  by  the  speed  with 

*  We  use  the  term  "apparent  activity"  for  the  activity  as  we  see  it.  In 
the  same  action  it  varies  with  the  conditions.  The  intrinsic  activity  or  affinity, 
on  the  other  hand,  is  the  absolute  activity  of  the  action  under  certain  carefully 
defined  conditions  (see  p.  295),  which  are  the  same  for  all  actions  to  be  compared. 


CHEMICAL  EQUILIBRIUM  291 

which  the  action  proceeds.  Thus,  if  the  foregoing  section  be  re- 
examined,  it  will  be  seen  that  we  spoke  throughout  of  the  speed, 
rather  than  of  the  tendency  or  activity. 

Finally,  temperature  and  other  conditions  influence  also  the 
activities  in,  and  therefore  the  speeds  of,  those  actions  which  pro- 
ceed to  completion,  and  are  not  reversible.  Hence,  unless  our 
statements  are  expressly  restricted  to  reversible  actions  and  to 
states  of  equilibrium,  they  apply  to  all  chemical  changes. 

The  Influence  of  Concentration.  —  In  the  first  place,  let  us 
assume  that  the  temperature  is  constant,  and  let  us  confine  our 
attention  for  the  present  to  the  influence  of  concentration  upon  a 
chemical  reaction.  We  have  seen  (p.  289)  that  the  speed  of  a  chemi- 
cal change  is  determined  by  the  frequency  with  which  the  molecules 
of  the  interacting  substances  encounter  one  another.  The  frequency 
of  the  encounters  amongst  a  given  set  of  molecules,  resulting  in  a 
definite  chemical  change,  will  in  turn  evidently  depend  entirely  upon 
the  degree  to  which  the  molecules  are  concentrated  in  each  other's 
neighborhood.  Larger  amounts  of  one  of  the  materials,  for  example, 
will  not  result  in  more  rapid  chemical  action  in  the  sense  which  this 
material  favors,  if  the  larger  amount  of  material  is  also  scattered 
through  a  larger  space.  Chemical  changes,  therefore,  are  not  acceler- 
ated by  increasing  the  mere  quantity  of  any  ingredient,  but  only  by 
increasing  the  concentration  of  its  molecules.  Thus,  a  large  amount 
of  hydrochloric  acid  with  a  piece  of  zinc  will  generate  hydrogen  no 
faster  than  a  smaller  amount.  But  substitution  of  more  concentrated 
acid  will  instantly  increase  the  speed  of  the  action.  In  the  second 
case,  the  number  of  molecules  of  the  acid  reaching  the  zinc  per 
second  is  greater,  and  this  action,  being  non-reversible,  proceeds 
more  rapidly  to  complete  consumption  of  the  zinc.  So  also,  iron 
burns  faster  in  oxygen  (100  per  cent)  than  in  air  (20  per  cent  oxygen). 

With  a  reversible  action  the  effect  on  the  speed  is  the  same,  ex- 
cepting that  the  continued  activity  of  the  reverse  action  prevents 
the  direct  one  from  reaching  completion.  Thus,  if,  in  the  action  of 
hydrogen  upon  iodine,  we  introduce  into  the  same  space  an  extra 
amount  of  hydrogen,  this  facilitates  the  formation  of  hydrogen  iodide 
by  increasing  the  possibilities  of  encounter  between  hydrogen  and 
iodine,  while  at  the  same  time  it  does  not  affect  (cf.  p.  Ill)  the  number 
of  encounters  in  a  given  time  between  hydrogen  iodide  molecules 
which  result  in  the  reverse  transformation.  The  proportion  of  hydro- 
gen iodide  formed,  therefore,  from  a  given  amount  of  iodine  will  be 


292  INORGANIC  CHEMISTRY 

greater,  although  the  total  possible  (by  complete  consumption  of  the 
materials)  has  not  been  altered,  since  the  quantity  of  one  ingredient 
only  has  been  increased.  The  introduction  of  an  excess  of  iodine 
would  have  had  precisely  the  same  effect. 

An  Experimental  Illustration.  —  It  is  easy  to  illustrate  this 
experimentally.  A  reaction  in  which  the  effects  of  different  concen- 
trations were  carefully  studied  by  Gladstone  (1855)  affords  a  good 
example.  If  ferric  chloride  and  ammonium  thiocyanate  are  mixed  in 
aqueous  solution,  a  liquid  containing  the  soluble,  blood-red  ferric 
thiocyanate  is  produced.  The  compound  radicals  are  NH4  and  CNS, 
and  the  action  is  a  simple  double  decomposition: 

FeCl3  +  3NH4CNS  <=±  Fe(CNS)3  +  3NH4C1. 

The  action  is  a  reversible  one,  and  the  system  is  homogeneous,  i.e., 
there  is  no  precipitation.  Now,  if  the  two  just-named  salts  are  mixed 
in  very  dilute  solution  in  the  proportions  required  by  the  equation, 
say  by  adding  20  c,c.  of  a  deci^bormal  solution  (p.  182)  of  each  to 
several  liters  of  water,  a  pale-r6ddish  solution  is  obtained.  When  this 
is  divided  into  four  parts,  and  one  is  kept  for  reference,  the  addition 
of  a  little  of  a  concentrated  solution  of  ferric  chloride  to  one  jar,  and 
of  ammonium  thiocyanate  to  another,  will  be  found  to  deepen  the 
color  by  producing  more  of  the  ferric  thiocyanate.  On  the  other 
hand,  mixing  a  few  drops  of  concentrated  ammonium  chloride  solution 
with  the  fourth  portion  will  be  found  to  remove  the  color  almost 
entirely  on  account  of  its  influence  in  accelerating  the  backward 
change. 

The  Law  of  Molecular  Concentration.  —  The  general  prin- 
ciple discussed  and  illustrated  in  this  section  may  be  called  the  law 
of  molecular  concentration,  and  may  be  stated  as  follows:  In  any 
given  chemical  change  the  apparent  activity,  and  therefore  the  speed 
of  that  change  is  proportional  to  the  molecular  concentration  of  each 
interacting  substance.  This  holds  whether  the  action  is  reversible 
or  not. 

Of  course,  when  different  actions  are  compared,  the  intrinsic 
affinities  will  be  different,  and  so,  with  equal  molecular  concentra- 
tions, the  speeds  will  be  different.  In  the  same  action,  a  change  in 
temperature  (see  p.  304)  will  alter  the  speed. 

The  phrase  "active  mass"  is  commonly  employed  instead  of  the  words 
"molecular  concentration."  It  is  distinctly  misleading,  however,  for.  as  we  have 


CHEMICAL  EQUILIBRIUM  293 

seen,  it  is  not  on  the  mass  of  a  substance,  but  on  the  quantity  of  it  in  a  given 
volume,  that  the  speed  of  the  action  depends.  If  a  physicist  spoke  of  the  mass, 
when  he  meant  density  (quantity  in  a  given  volume),  he  would  lose  all  scientific 
standing  at  once.  But  in  chemistry  it  is  not  considered  bad  form  to  use  the  word 
mass  in  this  connection,  although  a  conception  of  the  nature  of  density  is  intended. 
Some  kinds  of  nomenclature  are  used  more  loosely  in  chemistry  than  they  would 
be  in  any  other  science  (p.  210). 

The  following  is  a  more  exact  statement  of  the  law  of  molecular  concen- 
tration than  that  given  above:  The  speed  of  a  given  chemical  change  is  propor- 
tional to  the  first,  or  some  higher  power  of  the  molecular  concentration  of  each 
interacting  substance,  the  power  being  for  each  substance  determined  by  the 
number  of  molecules  of  that  substance  concurring  in  the  interaction  [or  by  the 
order  of  the  action  into  which  its  molecules  enter,  or  by  the  number  of  its  mole- 
cules appearing  in  a  molecular  equation  representing  the  action],  the  "action"  in 
question  being  the  slowest  of  the  partial  actions  involved,  when  the  action,  as  a 
whole,  is  achieved  through  consecutive  partial  actions  (see  p.  296). 

A  Warning.  —  The  reader  must  avoid  the  idea  that  a  reversible 
action  is  one  which  goes  to  completion,  and  then  runs  back  to  a 
certain  extent.  This  conception  would  be  contrary  to  the  fact,  and 
opposed  to  the  principles  of  energetics,  as  well  as  inexplicable  by  the 
kinetic-molecular  view. 

This  erroneous  idea  seems  to  assume  that  a  chemical  action  has 
momentum,  which  is  not  the  case.  If  a  reaction  did  have  momentum, 
and  ran  first  to  completion,  then,  on  reversal,  it  would  have  the  same 
momentum  and  would  come  back  to  the  starting  point  (the  original 
substances).  There  would  be  no  more  reason  for  its  stopping  at  an 
intermediate  position  on  the  return,  than  on  the  forward  journey. 
In  fact,  somewhat  like  a  pendulum  (which  does  have  momentum),  it 
would  swing  back  and  forth  and  never  come  to  rest,  or  rather  to  a 
state  of  balanced  activities!  A  pendulum,  when  pulled  out  of  the 
vertical  position,  starts  with  no  speed,  then  gains  a  little,  and  finally 
moves  with  its  greatest  velocity  when  passing  through  the  vertical 
position.  But  a  chemical  reaction  does  just  the  opposite.  The 
forward  action  starts  at  its  maximum  speed  (because  the  interacting 
molecules  are  at  a  maximum,  since  none  have  been  used  up)  and  goes 
more  and  more  slowly  until,  partly  because  the  interacting  molecules 
are  fewer,  and  partly  because  the  reverse  action  is  increasing  and 
undoing  its  work,  the  action  can  produce  no  further  change  and  the 
equilibrium  point  is  reached.  This  is  the  molecular  view  of  this 
fallacy. 

In  the  point  of  view  of  energy,  the  action  proceeds  at  first  because 
the  free  energy  in  the  system  is  diminishing.  At  the  equilibrium 


294  INORGANIC  CHEMISTRY 

point,  there  is  also  an  energy  equilibrium,  and  the  energy  of  the 
system  neither  diminishes  nor  increases.  To  carry  the  action  beyond 
the  equilibrium  point,  however,  would  require  an  increase,  from  some 
outside  source,  in  the  energy  of  the  system. 

As  to  the  facts,  a  system  involving  a  reversible  action  changes  in 
composition  while  approaching  equilibrium,  but  when  it  has  reached 
that  point,  no  further  changes,  such  as  the  pendulum  analogy  suggests 
can  be  detected.  To  be  specific,  if  the  ferric  chloride  and  ammonium 
thiocyanate  reaction,  described  above,  first  went  to  completion,  and 
then  receded  to  an  equilibrium  position,  the  mixture  would  first  be- 
come deep  red,  and  then  go  back  to  a  paler  color,  but  no  such  phenom- 
enon has  ever  been  observed. 

Formulation  of  the  Relation  between  Molecular  Concen- 
tration and  Speed  of  Reaction.  —  These  principles  will  become 
clearer  when  stated  in  somewhat  more  precise  terms.  We  confine  our 
attention  to  homogeneous  systems,  in  the  first  place. 

The  molecular  concentration  is  stated,  numerically,  for  each  sub- 
stance, in  terms  of  the  number  (whole  or  fractional)  of  moles  (gram- 
molecular  weights,  p.  236)  of  the  substance  contained  in  a  liter 
of  the  whole  mixture.  There  is  the  same  number  of  molecules  in  a 
mole  of  every  substance,  namely,  the  number  of  molecules  in  32  g.  of 
oxygen  (cf.  p.  238).  Hence  the  number  of  moles  per  liter  defines  the 
concentration  of  the  substances  in  terms  of  this  number  of  molecules 
in  a  liter  as  the  unit  of  concentration. 

Thus,  in  a  solution  containing  25.4  g.  of  free  iodine  (A  of  a  formula 
weight,  I2)  per  liter,  the  solution  is  0.1  molar  (p.  183),  and  the  molec- 
ular concentration  of  the  iodine  is  0.1.  When  the  substance  is  a  gas, 
the  concentration  of  the  molecules  is  proportional  to  the  partial  pres- 
sure of  the  gas.  Now,  one  mole  of  a  gas  occupies  22.4  liters  at  one 
atmosphere  pressure  (and  0°).  Hence,  when  one  mole  of  a  gas  is  con- 
tained in  1  liter,  and  its  molecular  concentration  is  therefore  1,  it  exer- 
cises 22.4  atmospheres  partial  pressure.  When  the  partial  pressure  of 
one  gas  in  a  mixture  is  two  atmospheres,  its  molecular  concentration  is 
ifr  or  0.09. 

The  speed  of  the  action  is  also  expressed,  numerically,  in  moles 
of  each  substance  transformed  per  minute  or,  for  slow  reactions,  per 
hour. 

Using  these  units,  the  relation  between  the  molar  concentrations 
and  the  speed  is  easily  expressed.  The  speed  is  proportional  to  the 
molar  concentration  of  each  molecule  appearing  in  the  molecular 


CHEMICAL  EQUILIBRIUM  295 

equation  for  the  action,  and  to  the  intrinsic  affinity  or  activity  of  the 
action.  For  example,  in  the  interaction  of  hydrogen  and  iodine 
(p.  287),  H2  +  I2  ->  2HI,  if  [H2]  and  [I2]  represent  the  molar  con- 
centrations, which  can  be  varied,  and  k  is  a  constant,  representing  the 
affinity,  which  is  invariable  for  each  action,  and  S  is  the  speed,  then: 

[H2]  X  [Ia]  X  ki  =  S. 

Similarly,  for  the  reverse  action:  2HI  — >  H2  +  I2: 
[HI]  X  [HI]  X  k2  =  [HI]2  X  k2  =  S. 

Note  that  [HI]  represents  the  total  concentration  of  HI  (not  one-half 
of  it),  because  every  molecule  of  HI  present  may  play  the  part  of 
either  one  of  the  two  demanded  by  the  equation. 

It  need  hardly  be  added  that,  clearly,  when  the  concentration  of 
any  ingredient  becomes  zero  (say  by  all  of  it  entering  into  combina- 
tion), or  when  the  affinity  is  zero,  the  speed  S  must  become  zero,  that 
is  to  say,  no  action  takes  place. 

The  affinity  or  activity  (k)  is  always  represented  numerically  by 
the  speed  in  moles  per  unit  of  time  when  the  moler  concentration  of 
each  ingredient  is  unity  (see  next  section). 

More  Specific  Formulation  of  Speed  of  Reaction.  —  In  the 

preceding  section  attention  was  not  called  to  one  practical  difficulty.  Since,  as 
the  action  proceeds,  the  materials  are  being  used  up,  their  concentrations  do  not 
remain  constant,  but  progressively  diminish  during  the  period  of  observation  (1 
minute  or  1  hour) .  Thus,  the  action  becomes,  on  this  account,  slower  and  slower. 
S,  therefore,  as  used  above,  is  the  speed  at  a  given  moment.  If  applied  to  one 
minute  (or  one  hour),  it  is  the  amount  which  would  be  transformed  if  the  con- 
centrations present  at  that  moment  were  artificially  maintained  during  the  whole 
minute  (or  hour). 

Now,  the  theoretical  speed  at  constant  concentration,  used  in  the  above 
formulae,  can  be  calculated  from  the  speeds  observed  with  diminishing  concen- 
trations. It  will  be  noted  that  when  the  concentrations  are  unity,  [H2]  =  [la]  =  1, 
and  therefore  1  X  1  X.  k  =  S.  That  is,  with  unit  concentrations,  k  =  S.  The 
plan,  therefore,  is  to  experiment  with  different  concentrations  at  the  same  tem- 
perature, and  to  measure  the  amounts  transformed  in  measured  times  in  each 
case.  We  can  then  calculate  (see  below)  from  the  data  of  each  set  the  speed 
(moles  transformed  per  hour  or  minute)  which  would  be  shown  by  constantly 
maintained  unit  concentrations  of  the  materials.  The  answers  are  the  values  of 
k  for  unit  concentrations,  based  upon  measurements  with  different  concentrations 
of  the  same  substances,  and  they  agree  closely  with  one  another.  The  values  of  k 
for  several  different  chemical  actions,  however,  may  differ  widely,  and  are  meas- 
ures of  the  relative  activity  of  each. 


296  INORGANIC  CHEMISTRY 

The  relation  of  the  theoretical  speed  with  constant  unit  concentration  to  that 
which  is  observed  with  diminishing  concentration  is  as  follows:  For  an  action  of 
the  form:  A  —  >  B  +  C,  where  the  change  in  only  one  molecule  constitutes  the 
action,  if  [A]  is  the  initial  molecular  concentration  of  A,  and  x  is  the  fraction  of 
this  which  is  transformed  in  the  time  t, 

k  =  $unit  concn.  =  loge  r"    ,  _      4-  t. 

When  two  molecules  have  to  interact:  A  +  B  —  »  C  +  D,  the  formula  is  still  more 
complex.  If  the  substances  are  present  in  equivalent  proportions,  their  molec- 
ular concentrations  in'  this  special  case  are  alike,  and  may  each  be  represented 
by  [A],  for  [A]  =  [B].  The  relation  is  then: 


The  mathematical  derivation  of  these  relations  will  be  found  in  any  work  on 
physical  chemistry. 

An  Illustration.  —  The  following  illustration  (see  also  Sulphurous  acid) 
will  make  all  this  clearer.  When  arsine  AsH3  (q.v.}  is  heated  at  310°,  it  decom- 
poses gradually  into  hydrogen  and  arsenic: 

AsH3-»As 


The  action  is  not  appreciably  reversible.  The  arsenic  assumes  the  solid  form. 
The  gas  is  inclosed  in  a  tube  which  is  kept  in  a  bath  at  310°,  and  a  manometer 
shows  changes  in  pressure.  Since,  as  the  action  proceeds,  1|  molecules  of  hydro- 
gen take  the  place  of  each  molecule  of  arsine,  the  total  pressure  slowly  increases. 
Every  increase  of  1  mm.  in  the  pressure  is  the  result,  therefore,  of  an  addition  of 
3  mm.  partial  pressure  of  hydrogen  and  a  reduction  of  2  mm.  in  the  partial  pres- 
sure of  the  arsine.  The  molecular  concentrations  are  proportional  to  the  pres- 
sures, and  change,  therefore,  in  the  same  ratios.  The  observations  (first  two 

*  Ordinarily  we  should  write  the  equation  2AsH3  —  »  2As  +  3H2.  But  this 
form  would  make  the  speed  proportional  to  [AsH3]2  and  calculation  would  then 
give  us  inconstant  values  for  S.  The  reason  for  this  apparent  irregularity  is  that 
the  action  takes  place  in  two  stages  (consecutive  reactions,  p.  272).  First  AsHs 
decomposes  to  give  As  +  3H.  Then  the  atoms  of  hydrogen  combine  in  pairs  to 
form  H2.  The  latter  of  these  actions,  however,  is  very  speedy,  while  the  former 
is  the  one  that  takes  time,  and  it  is  the  time  occupied  by  the  former  alone  that  we 
are  measuring.  Hence,  only  the  former  is  involved  in  the  calculation.  This 
point  is  covered  by  the  last  part  of  the  more  exact  statement  of  the  law  of  molec- 
ular concentration  (p.  293).  The  rates  of  very  fast  actions  cannot  be  measured. 
In  a  sense,  it  is  the  slowness  of  the  action,  that  is,  the  time  required  for  a  certain 
amount  of  chemical  change,  that  we  are  measuring  when  we  determine  so-called 
speeds  of  reaction. 


CHEMICAL  EQUILIBRIUM 


297 


columns),  together  with  the  data  deduced  from  the  first  two  by  calculation,  were 
as  follows: 


MOLECULAR  CONCENTRATIONS. 

Total. 

AsH3  Tranafmd. 

<Sunit  conon.  =  k. 

0 

784.84 

0.02159 

3 

878.50 

0.02416 

6.00514 

0.0906 

4 

904.05 

. 

8 

987.19 

.... 

We  must  first  ascertain  the  molecular  concentration  of  the  arsine  correspond- 
ing to  the  observed  pressure  at  the  beginning.  We  remember  that  at  22.4  atmos- 
pheres, or  22.4  X  760  mm.  and  0°,  the  molar  concentration  of  a  gas  has  the 
value  1  (p.  294).  The  actual  initial  pressure  784.84  mm.  at  310°  would  become. 

784  84  X  273 

at  0°, '• ,  or  367.5  mm.    The  molecular  concentration  is  here,  therefore, 

(olO  -p  27 o) 

or  0.02159  moles  per  liter.     After  3  hours,  some  hydrogen  has  been 

22.4  X  760 

formed.  The  pressure  has  increased  to  878.50  mm.  Reducing  as  before,  this 
represents  a  molecular  concentration  of  all  ingredients  of  0.02416  moles  per  liter. 
The  increase  is  0.00257.  This,  as  was  demonstrated  above,  corresponds  to  a  loss 
of  2  X  0.00257,  or  0.00514  moles  per  liter  of  arsine.  Now,  employing  the  formula 
given  above,  we  find  the  speed  per  hour: 


Smut  concn.   = 


[AsH3]  -  x 


0.01645 


This  result  means  that,  if  the  concentration  of  the  arsine  were  to  be  maintained 
at  the  initial  value  by  continual  renewal  of  the  waste,  then  0.0906  (9.06  per  cent) 
of  the  initial  amount  would  be  decomposed  in  an  hour.  Using  the  pressures  at 
4  and  at  8  hours,  the  reader  will  obtain  by  calculation,  practically  the  same  value 
for  k.  Other  experiments  with  still  different  concentrations,  provided  the  tem- 
perature was  the  same  (p.  291),  would  likewise  give  the  same  results,  for  this  is 
the  speed  calculated  back  to  unit  concentration.  Thus  the  affinity  or  activity 
of  the  action  may  be  measured  with  any  concentration,  and  expressed  in  moles 
transformed  per  hour  with  unit  concentration.  Similar  measurements  with  other 
actions  then  enable  comparisons  of  their  relative  activities  to  be  made  (see  Exer- 
cise 10,  end  of  this  chapter). 

Formulation  of  the  Conditions  for  Chemical  Equilibrium. 

—  The  plan  outlined  above  (p.  295)  may  be  used  further  to  formulate 
the  conditions  for  chemical  equilibrium.  As  we  have  seen  (p.  289), 


298  INORGANIC  CHEMISTRY 

a  characteristic  of  a  system  in  chemical  equilibrium  is  that  the  speeds 
of  the  forward  and  reverse  actions  have  become  equal.  If,  then, 
[H2]eqm.,  [Weqm.,  and  [HI]eqm.,  now  represent  the  molecular  concen- 
trations when  the  system  has  reached  equilibrium,  then,  since  the 
speeds  are  equal: 

[H2]eqm.  X   [I2]eqm.  X  ki  =   [Hl]2eqm.  X  k2 

and 

PBtUn.  X  [I2]eqm.  =  k2  =  congtant  =  K 
[HI]2eqm.  ki 

In  words,  this  means  that  if  we  change  the  amount  of  hydrogen  iodide 
placed  in  the  same  vessel,  or  if  we  use  amounts  of  hydrogen  and  iodine 
which  are  not  equivalent,  the  numerical  value  at  equilibrium  of  each 
concentration  ([H2],  etc.),  will,  of  course,  be  different,  but  the  product 
of  the  concentrations  of  hydrogen  and  of  iodine,  divided  by  the 
concentration  of  the  hydrogen  iodide,  will  always  give  the  same 
numerical  value  for  the  constant  at  the  same  temperature.  This 
numerical  value  is  called  the  equilibrium  constant. 

If,  for  example,  the  value  of  the  constant  is  J,  then  the  speeds  of 
the  two  actions,  if  each  were  to  proceed  unimpeded  (say  in  separate 
vessels)  with  constantly  maintained  unit  concentrations  of  the 
materials,  would  be  in  the  ratio  k2  :  ki  or  1  :  4.  From  this  it  will  be 
seen  that  measurement  of  the  concentrations  present  in  a  system 
which  has  reached  equilibrium  gives  us  the  data  for  calculating  the 
value  of  this  ratio.  In  other  words,  it  gives  us  the  means  of  ascertain- 
ing the  relative  magnitudes  of  the  intrinsic  affinities  of  the  opposed 
actions. 

Applying  this  to  the  data  given  (p.  290)  for  hydrogen  iodide, 
at  445°,  with  equivalent  quantities  of  the  two  elements,  nearly  0.8 
(more  exactly,  79  per  cent)  of  the  weight  of  each  is  in  the  form  HI, 
and  0.2  in  the  mixture  H2  +  I2.  Thus  in  every  100  molecules,  80  are 
HI,  10  are  H2,  and  10  are  I2.  Thus  K  =  O.I2  ^  0.82  =  ^.  That  is 
to  say,  the  union  of  hydrogen  and  iodine  would  take  place  with  64 
times  as  great  a  speed  as  the  dissociation  of  hydrogen  iodide  if  each 
action  could  proceed  without  reversal  and  under  identical  conditions. 
Or,  in  terms  of  the  kinetic  theory,  the  collisions  of  the  H2  and  I2 
molecules  result  many  times  more  often  in  chemical  change  than  do 
collisions  of  HI  molecules. 

The  case  of  hydrogen  iodide  is  comparatively  simple  because  the 
volume  is  not  altered  by  the  progress  of  the  action  (see  below).  The 


CHEMICAL  EQUILIBRIUM  299 

expansion  when  phosphorus  pentachloride  (p.  260)  dissociates  compels 
us  to  take  account  of  the  volume.     The  equation  is: 

Cla         and       K 


If  one  gram  molecule  of  the  pure  PCls  is  taken,  and  x  is  the  proportion 
dissociated,  and  v  the  volume  occupied  by  the  whole,  then  [PCl3]eqm.  = 

[CMeqm.    "f  ^>  ^d   [PClJeqm.  =    ~  —  .       Thus   K   =          ^  NOW 

at  250°  (and  760  mm.),  for  example,  0.8  of  the  whole  weight  of  mate- 
rial is  dissociated  :  x  =  0.8  and  1  -  x  =  0.2.  Hence  K  =  0.82  -f-  Q.2v 
=  3.2  -f-  v.  To  obtain  the  value  of  v  we  note  that  a  gram  molecule 
at  760  mm.  and  0°  occupies  22.4  liters.  At  250°  it  occupies  22.4  X 
(250  +  273)  -5-  273  1.  But  this  mass  of  gas  contains  0.8  more  mole- 
cules because  of  dissociation,  and  its  volume  is,  therefore,  1.8  X 
22.4  (250  +  273)  -5-  273  =  v  =  77.2  1.  Thus  K  =  3.2  -r-  77.2  =  ^. 
Otherwise  stated,  the  union  of  the  trichloride  and  chlorine  would 
proceed  twenty-five  times  as  fast  as  the  dissociation,  if  each  of  the 
three  substances  was  present  in  unit  concentration,  and  each  action 
could  proceed  independently  without  reversal. 

The  Effect  of  Changes  of  Volume  on  Chemical  Equilibrium. 

—  Our  applications  of  the  theory  of  equilibrium  will  be  chiefly  to  dis- 
solved bodies,  and  hence  the  effect  on  the  equilibrium  point  of  changes 
in  volume  (by  dilution  or  the  reverse)  will  require  frequent  considera- 
tion. Now  dilution,  for  example,  diminishes  opportunities  for  en- 
counters between  the  substances  on  both  sides  of  the  equation.  In  the 
second  of  the  above  illustrations,  increasing  the  volume  decreases  the 
rate  at  which  the  chlorine  and  the  phosphorus  trichloride  can  com- 
bine. Since,  however,  the  speed  of  the  dissociation  depends  on  the 
state  of  the  PCU  molecules  only,  and  is  unaffected  by  their  nearness  to 
or  remoteness  from  one  another,  the  forward  action  will  not  be 
weakened  at  all.  Hence,  dilution  increases  the  degree  of  dissociation. 
In  general,  change  in  volume  will  affect  the  equilibrium  point  when- 
ever there  are  more  molecules  on  one  side  of  the  equation  than  on  the 
other. 

In  mathematical  terms,  when  we  change  the  volume  to  -  times  its 

n 

former  value  (n  whole  or  fractional),  the  concentration  changes  n 
times.  The  equation  for  equilibrium  then  becomes,  momentarily, 


300  INORGANIC  CHEMISTRY 

n  [PC13]  X  n  [C12]  -s-  n  [PC15]  *  K,  orn  [PC13]  X  [C12]  -*-  [PCU]  ^  K. 
To  restore  the  value  of  the  expression  to  equality  with  K,  change  must 
occur  in  the  concentrations  [PC13],  [C12],  and  [PCls].  When  n  is  <  1, 
that  is,  when  the  volume  increases,  some  PCls  must  pass  into  the  form 
PC13  and  C12  until  [PCls]  X  [C12]  -  [PCU1  =  K,  as  before.  That  is, 
dilution  increases  the  degree  of  dissociation. 

In  case  of  hydrogen  iodide,  and  in  all  others  where  the  number  of  molecules 
taking  part  in  the  direct  and  reverse  actions  is  the  same,  change  in  the  volume  of 
the  system  has  no  effect  on  the  position  of  the  equilibrium  point.  Thus  dilution 
diminishes  the  chance  of  encounter  between  two  HI  molecules  to  the  same  extent 
that  it  interferes  with  encounters  between  H2  and  I2  molecules.  Conversely,  in- 
crease in  all  concentrations,  by  diminution  of  volume,  favors  both  actions  equally. 
Hence,  at  445°,  79  per  cent  of  HI  will  always  be  present  at  last,  whatever  the 
volume  occupied  by  a  given  amount  of  the  materials.  In  mathematical  terms,  if 
we  diminish  the  volume  n  times  (n  whole  or  fractional),  we  increase  the  concen- 
tration of  each  constituent  n  times.  The  values  become  n  [H2],  n  [I2],  and  n  [HI] 
respectively: 

n'[Hj  XIW      [H,l  X[I«1 


and  K  = 


n2  [Hip  [Hip 


Heterogeneous  Equilibrium.  —  A  modification  of  the  above 
conceptions  is  necessary  when  the  mixture  is  not  homogeneous.  If, 
for  example,  one  of  the  constituents  is  present  as  a  solid  or  a  gas,  in 
greater  amount  than  can  be  dissolved  by  the  liquid  in  which  alone  the 
chemical  change  takes  place,  then,  according  to  the  definition  of 
saturated  solution  (p.  192),  the  concentration  of  the  dissolved  material 
will  be  constant  at  a  given  temperature  as  long  as  physical  equilibrium 
between  the  solid  and  the  solution  is  maintained.  This  is  a  case 
especially  likely  to  occur  when  slightly  soluble  (so-called  "insoluble") 
bodies  (cf.  p.  179)  are  concerned. 

The  same  reasoning  applies  also  to  very  slightly  volatile  solids. 
The  concentration  of  the  vapor  of  a  solid  body  present  in  excess  (meas- 
ured by  its  vapor  pressure)  will  be  constant  so  long  as  the  temperature 
is  fixed,  and  interaction  with  a  superincumbent  gas  will  take  place 
chiefly  through  the  vapor. 

In  both  these  cases  the  concentrations  of  the  active  parts  of  the 
slightly  soluble  and  slightly  volatile  bodies,  respectively,  are  not  sub- 
ject to  variation  —  they  are  constant.  Thus,  with  the  action, 

CuSO4,H20  -»  CuS04  +  H20  T 
the  concentrations  of  the  vapors  [CuS04,H20]  and  [CuSO4]  are  con- 


CHEMICAL  EQUILIBRIUM  301 

slant,  and  utterly  negligible  and  that  of  water  [H20]  alone  is  subject 
to  alteration.  We  have,  therefore, 

[CuS04]  X  [H20]  [CuS04,H20]T^ 

[CuS04,H20]  [CuS04]    K 

in  which,  since  [CuSO4,H2O],  [CuSO4],  and  K  are  constant,  [H2O]  must 
be  constant  also.  But  the  pressure  of  a  gas  is  proportional  to  its 
molecular  concentration,  according  to  Avogadro's  hypothesis.  There- 
fore, in  this  action,  the  pressure  of  the  water  vapor  (the  dissociation 
pressure)  should  be  constant  irrespective  of  the  extent  to  which  the 
dissociation  has  progressed.  Observation  shows  that  this  is  the  case. 
All  hydrates,  as  we  have  seen  (pp.  151-154),  behave  in  a  precisely 
similar  way,  and  furnish  numberless  confirmations  of  this  application 
of  the  law  of  molecular  concentration. 

Applications:  The  Reverse  Action.  Displacement  of  Equi- 
libria. —  Of  special  interest  to  the  chemist  are  the  conditions  under 
which  the  equilibrium  point  may  be  displaced  and  more  nearly 
complete  realization  of  one  of  the  two  opposed  changes  may  be 
brought  about. 

We  have  seen  (p.  292)  that  one  way  in  which  a  reversible  action 
may  be  forced  nearer  to  completion  in  one  direction  or  the  other  is  the 
introduction  of  an  excess  of  one  of  the  ingredients  contributing  to  the 
action.  This  method  of  displacing  the  equilibrium  point,  however, 
cannot  be  very  effective  unless  it  is  possible  to  introduce  an  exceed- 
ingly large  excess  of  the  selected  ingredient  in  a  high  degree  of  molec- 
ular concentration,  since  this  operation  does  not  in  any  way  effect  or,  in 
particular,  restrain  the  reverse  action  which  is  continually  undoing  the 
work  of  the  forward  one.  A  much  more  effective  means  of  furthering 
the  desired  direction  of  such  actions  is  found,  therefore,  in  the  restraint 
or  practical  annulment  of  the  reverse  action.  A  good  way  to  accom- 
plish this  is  to  allow  the  products  of  the  direct  action  to  separate 
into  an  inhomogeneous  mixture.  Any  agency  which  could  remove 
the  free  iodine  as  fast  as  it  was  formed  in  the  decomposition  of 
hydrogen  iodide,  for  example,  would  entirely  stop  the  reproduction 
of  the  compound  and  so  would  enable  the  dissociation  2HI  ^±  H2  +  I2 
to  run  to  completion.  The  concentration  of  one  product  can  often 
be  reduced  practically  to  zero.  To  achieve  the  same  effect  by  adding 
an  interacting  substance,  the  concentration  of  the  latter  would  have 
to  be  raised  to  infinity,  which  is  impossible. 


302  INORGANIC  CHEMISTRY 

This  might  be  realized  *  by  causing  one  end  of  a  sealed  tube 
charged  with  hydrogen  and  iodine,  after  the  contents  had  settled  down 
to  a  condition  of  equilibrium,  to  project  from  the  bath  in  which  the 
whole  had  been  kept  at  445°  (Fig.  97,  which  is  simply  diagrammatic). 

By  cooling  this  end,  a  large  part  of 
the  21  per  cent  of  free  iodine  would 
quickly  be  condensed  in  it  to  the  solid 
form,  while  the  hydrogen  would  re- 
main gaseous.  Only  the  trace  of 
vapor  which  cold  iooline  gives  would 
then  be  available  to  interact  with  the 

hydrogen  and  reproduce  hydrogen  iodide.  Meanwhile  the  decom- 
position of  the  latter  would  go  on,  and  thus,  eventually,  almost  all 
the  iodine  would  be  found  free  in  one  end  of  the  tube,  and  the 
hydrogen,  all  free  likewise,  would  occupy  the  rest.  By  this  purely 
mechanical  adjustment  the  chemical  change  would  in  this  way  be 
carried  from  21  per  cent  completion  to  almost  absolute  completion: 

2HI  *=*  H2  +  I2  (vapor)  *=»  I2  (solid). 

If,  on  the  other  hand,  arrangements  were  made  to  have  powdered 
marble,  in  a  sealed  bulb  of  thin  glass,  enclosed  in  the  tube,  we  might 
imagine  the  very  opposite  effect  of  the  above  to  be  produced.  The 
breaking  of  the  bulb  of  marble,  when  equilibrium  had  been  reached, 
would  provide  means  for  the  removal  of  all  the  hydrogen  iodide,f 
while  the  hydrogen  and  iodine  would  still  be  gaseous.  Thus,  the  com- 
pound having  been  removed,  there  would  be  no  reverse  action  to 
compensate  for  the  union  of  the  elements.  The  whole  material  would, 
therefore,  soon  have  passed  through  the  form  HI.  Hence,  by  another 
mechanical  arrangement,  an  action  which  ordinarily  could  progress  to 
only  79  per  cent  would  be  turned  into  a  complete  one. 

The  discussion  of  hydrogen  iodide  in  this  chapter  shows  very 
clearly  why  we  do  not  prepare  the  compound  by  uniting  the  elements 
(p.  277).  (1)  Since  the  elements  interact  as  gases,  very  bulky  appara- 
tus would  be  required  to  prepare  any  considerable  quantity;  (2)  the 

*  For  another  illustration,  see  under  Ammonia. 

t  The  hydrogen  iodide  would  be  destroyed  by  interaction  with  the  marble: 

2HI  +  CaCO3  — >  CaI2  +  CO2  +  H2O. 

The  calcium  iodide  is  a  solid.  The  two  gases,  carbon  dioxide  and  water  vapor, 
do  not  interact  with  hydrogen  or  with  iodine,  and  would  not,  therefore,  inter- 
fere with  the  formation  of  fresh  hydrogen  iodide, 


CHEMICAL  EQUILIBRIUM  303 

union  is  very  slow,  taking  many  hours  at  283°  to  give  82  per  cent; 
(3)  it  is  incomplete,  at  best,  and  we  obtain  a  mixture,  and  not  a  pure 
substance. 

Reversibility  Usually  Avoided.  —  In  every-day  chemical  work, 
since  our  object  is  usually  to  prepare  some  one  substance,  chemists 
either  avoid  chemical  changes  which  are  notably  reversible,  or  adjust 
the  conditions,  as  is  done  in  the  foregoing  illustrations,  so  that  the 
reverse  of  the  action  which  they  desire  is  prevented.  In  consequence 
of  this,  when  carrying  out  the  directions  for  making  familiar  prepara- 
tions, the  fact  that  such  actions  are  reversible  at  all  very  readily 
escapes  our  notice.  Arranging  the  conditions  so  that  the  separation 
of  a  solid  body  by  precipitation,  or  the  liberation  of  a  gas,  takes  place, 
are  the  two  commonest  ways  of  rendering  a  reversible  action  com- 
plete. Excellent  examples  of  both  of  these  are  furnished  by  the 
chemical  change  used  in  producing  hydrogen  chloride  by  the  inter- 
action of  salt  and  sulphuric  acid,  the  discussion  of  which  (p.  208) 
should  now  be  studied  attentively  in  the  light  of  these  explanations. 

The  escape  of  one  member  of  a  system  engaged  in  chemical  interaction, 
because  it  is  gaseous  or  solid,  and  in  either  case  immiscible  with  the  rest  of  the 
members  of  the  system,  is  the  commonest  cause  of  the  obstruction  of  one  direc- 
tion of  a  reversible  action  and  the  triumph  of  the  other.  This,  as  we  have  seen, 
is  the  combined  result  of  the  natural  behavior  of  a  system  in  chemical  equilibrium, 
and  of  the  physical  properties,  particularly  the  solubility,  of  the  members  of  the 
system.  Two  rules,  attributed  to  Berthollet,  have  been  made,  however,  to  de- 
scribe these  special  cases  of  a  broader  principle.  Unfortunately,  it  is  difficult  so 
to  word  them  that  they  shall  be  entirely  unambiguous  and  entirely  correct. 

The  "rule  of  precipitation,"  for  example,  might  read:  When  certain  classes  of 
materials  are  brought  together  in  solution,  if  an  exchange  of  radicals  would  produce 
an  insoluble  body,  this  exchange  will  occur.  But  then  the  fact  is  that,  in  such 
cases,  the  exchange  always  occurs  to  some  extent  whether  any  product  is  insoluble 
or  not.  The  insolubility  is  responsible  only  for  the  greater  completeness  of  the 
exchange.  Crude  statements  to  the  effect  that  "when  an  insoluble  body  can 
be  formed,  it  will  be  formed,"  when  close  scrutiny  shows  them  to  possess  any 
definite  meaning  whatever,  are  grossly  misleading.  They  suggest  that  insolubility 
is  a  sort  of  especially  desirable  career  on  which  the  elements  are  ambitious  of 
entering. 

All  forms  of  these  so-called  laws  are  objectionable,  because  they  necessarily 
suggest  that  the  positive  direction  of  the  action  is  assisted  by  the  immiscibility  of 
the  product,  and  this  is  the  precise  converse  of  the  fact.  The  immiscibility  does 
nothing  at  all  towards  assisting  the  formation  of  the  insoluble  substance  itself,  but 
does  whatever  it  can  towards  preventing  the  destruction  of  that  substance,  once  it 
is  formed,  by  hampering  the  negative  action. 


304  INORGANIC  CHEMISTRY 

Affinity  vs.  Solubility.  —  The  question  of  the  relation  of  affin- 
ity to  the  apparently  much  greater  efficiency  of  one  of  the  directions 
of  some  reversible  actions,  may  now  be  put  in  a  much  clearer  light  (pp. 
209,  127,  and  this  Chap.).  The  whole  of  the  possibilities  of  progress 
for  any  action  are  expressed  by  a  function  (p.  295)  of  the  form  dc^k  = 
S.  If  any  one  of  the  variables,  say  one  of  the  concentrations  (ci),  is 
negligible,  the  product  must  be  small,  irrespective  of  the  values  of  the 
other  factors.  Thus  the  feebleness  of  a  chemical  action  only  shows 
that  the  product  of  all  the  variables  is  minute,  and  not  that  the 
affinity  factor  per  se  is  of  small  magnitude. 

History.  —  The  conceptions  discussed  in  this  chapter  are  not 
new,  although  they  have  come  into  general  use  rather  recently.  The 
law  of  reaction  speed,  and  the  influence  of  the  concentrations  of  the 
reacting  substance  thereon  (p.  291),  was  set  forth  and  formulated  by 
Wilhelmy  as  early  as  1850.  Gladstone  (1855)  studied  quantitatively 
the  influence  of  concentration  in  cases  of  chemical  equilibrium  (p. 
292).  The  kinetic  explanation  (p.  289)  was  developed  by  Williamson 
(1851).  Finally  the  laws  of  chemical  equilibrium  were  formulated 
more  explicitly  and  applied  more  thoroughly  by  two  Norwegian 
chemists,  Guldberg  and  Waage  (1864-9). 

The  Influence  of  Temperature  on  the  Speed  of  any  Reac- 
tion. —  The  activity  of  chemical  change,  and  therefore  the  speed  of 
all  chemical  changes,  is  increased  by  raising  the  temperature  and 

diminished  by  lowering  it  (cf.  p.  93).  Thus,  zinc  displaces  hydrogen 
more  rapidly  from  hot  than  from  cold  hydrochloric  acid.  Different 
actions  are  affected  in  different  degrees,  and  no  simple  rule  accu- 
rately defining  the  effect  can  be  given.  Roughly  speaking,  however, 
a  rise  of  10°  doubles  the  speed  of  every  action.  A  rise  of  100°  will 
therefore  make  the  speed  roughly  1024  times  greater.  Hence,  when 
the  chemist  finds  that  two  substances  show  no  evidence  of  interac- 
tion, he  infers  that  there  must  be  either  slow  action  or  none,  and  he 
seeks  to  settle  the  question  quickly  by  heating  the  mixture. 

The  Influence  of  Temperature  on  a  System  in  Equilib- 
rium. —  In  a  reversible  change  the  two  opposing  reactions  are 
different  actions  and  their  speeds  are  therefore  affected  in  different 
degrees  by  the  same  alteration  in  temperature.  Hence,  when  the 
temperature  is  changed,  the  relative  amount  of  the  two  sets  of 
materials  present  is  altered  and  the  equilibrium  is  displaced.  Thus, 


CHEMICAL  EQUILIBRIUM  305 

in  Deacon's  process,  a  rise  of  40°  in  the  temperature  displaces  the 
equilibrium  backward  (p.  217),  and  diminishes  the  yield  of  chlo- 
rine by  5  per  cent.  In  the  vapor  of  phosphorus  pentachloride 
(p.  260),  the  displacement  is  in  the  opposite  direction.  The  vapor 
is  a  mixture  of  the  pentachloride  with  the  trichloride  and  free  chlorine : 
PC15  +±  PC13  +  C12.  At  200°  and  760  "mm.,  51.5  per  cent  of  the 
material  is  present  as  pentachloride  and  48.5  per  cent  as  trichloride 
and  chlorine.  Raising  the  temperature  to  250°  (760  mm.)  changes  the 
proportions  to  20  per  cent  and  80  per  cent,  respectively.  At  300° 
only  3  per  cent  of  the  pentachloride  remains.  Evidently,  here,  raising 
the  temperature  favors  the  decomposition  of  the  pentachloride, 
and  therefore  increases  the  speed  of  its  dissociation  more  than  it 
does  the  speed  of  the  reunion  of  the  trichloride  and  chlorine. 

Fcm't  Hoff's  Law.  —  One  use  of  a  law  is  to  enable  us  to  answer 
a  question,  when  we  have  not  in  memory  the  fact  constituting  the 
answer,  and  even  when  we  have  never  read  or  heard  the  fact.  The 
law  or  rule  enables  a  little  reasoning  to  take  the  place  of  a  vast 
amount  of  memorizing.  Thus,  to  answer  the  question:  Does 
sodium  chloride  always  have  the  same  composition,  it  is  not  necessary 
to  have  read  and  to  remember  all,  or  any  of  the  numerous  investiga- 
tions of  this  substance  that  have  been  made.  We  simply  refer  the 
question,  mentally,  to  the  law  of  definite  proportions,  and  say  "yes." 
Now  the  facts  mentioned  above  are  connected  by  a  law  which  will 
answer  many  practical  questions  in  chemistry. 

When  phosphorus  trichloride  and  chlorine  combine  (to  form 
PCls),  heat  is  given  out.  Conversely,  when  phosphorus  penta- 
chloride dissociates,  heat  is  absorbed:  * 

PC15  +  30,000  cal.  <*  PC13  +  C12. 

Now,  when  the  temperature  is  raised,  the  action  proceeds  in  the 
direction  of  decomposing  more  of  the  pentachloride.  That  is,  the 
equilibrium  is  displaced  in  the  direction  which  absorbs  heat. 

In  Deacon's  process,  we  find  that  the  interaction  of  hydrogen 
chloride  and  oxygen  liberates  heat, 

4HC1  +  O2  ?±  2H2O  +  2C12  +  28,000  cal., 

and  in  this  action  raising  the  temperature  drives  the  equilibrium 
backward,  and  a  lowering  in  the  temperature  is  required  to  increase 
the  yield  of  chlorine. 

The  rule  is  obvious,  and  applies  to  all  reversible  reactions :  When 


306  INORGANIC  CHEMISTRY 

the  temperature  of  a  system  in  equilibrium  is  raised,  the  equilibrium 
point  is  displaced  in  the  direction  which  absorbs  heat.  In  other 
words,  a  rise  in  temperature  favors  the  interaction  of  that  one  of 
the  two  sets  of  materials  to  which  the  heat  is  added  (+  sign)  in  the 
equation.  If  the  equation  happens  to  be  written  with  a  negative 
heat  of  reaction  (e.g.,  p.  100),  the  heat  can,  of  course,  be  transferred 
to  the  other  side  with  its  sign  changed.  This  law  is  'known  as  van't 
Hoff's  law  of  mobile  equilibrium. 

We  have  already  encountered  numerous  illustrations  of  this  law. 
The  interaction  of  steam  and  iron  (p.  116)  is  exothermal,  and  so  the 
higher  the  temperature,  the  more  conspicuous  the  reverse  action 
becomes.  Again,  as  the  temperature  rises,  barium  peroxide  gives  a 
higher  pressure  of  oxygen  (p.  82),  hydrates  give  a  greater  pressure  of 
water  vapor  (p.  151),  and  the  dissociation  of  molecular  hydrogen 
increases  (p.  253),  because  these  actions  all  absorb  heat.  Many  other 
examples  will  be  noted  as  we  reach  them  (see  ozone,  ammonia,  nitric 
oxide). 

This  law  is  of  practical  value.  More  than  once,  in  chemical 
factories,  much  time  and  money  have  been  spent  on  trying  to  arrange 
machinery  to  give  a  better  yield  of  some  substance  at  a  high  tem- 
perature, when  a  reference  to  this  law  would  have  shown  that  the 
chief  change  necessary  was  to  use  a  lower  temperature,  and  perhaps 
hasten  the  action  by  use  of  a  contact  agent. 

Application  to  Hydrogen  Iodide.  —  At  283°,  a  mixture  of 
hydrogen  and  iodine  yields  82  per  cent  of  hydrogen  iodide  and  18 
per  cent  of  the  uncombined  elements.  At  445°,  the  yield  of  hydrogen 
iodide  is  79  per  cent,  and  at  508°  only  76  per  cent.  Since  the  elements 
increase  in  quantity  as  the^'temperature  rises,  we  infer  that  the 
dissociation  of  the  compound  absorbs  heat.  At  400°,  the  value  is: 

2HI  +  535  cal.  -»  H2  +  I2. 

Curiously  enough,  at  low  temperatures,  the  action  is  exothermal. 
Thus  at  18°: 

2HI->H2  +  I2  +  6100  cal. 

A  reversal  of  the  sign  of  the  heat  of  a  reaction  is  not  uncommon. 
Thus,  ammonia  and  hydrogen  bromide,  up  to  about  320°,  give  out 
heat  in  combining.  Beyond  that  temperature,  ammonium  bromide 
gives  out  heat  in  dissociating  (A.  Smith),  and  so  beyond  that  tem- 
perature the  degree  of  dissociation  is  less  the  higher  the  temperature 


CHEMICAL  EQUILIBRIUM  307 

(with  135  mm.  pres.  35  per  cent  at  330°,  and  20  per  cent  at  390°). 
Such  reversals  are  common  in  cases  of  ionization  (q.v.). 

Application  to  Physical  Equilibria.  —  Van't  Hoff's  law  ap- 
plies also  to  physical  processes.  Thus,  as  the  temperature  rises, 
a  substance  which  absorbs  heat  in  dissolving  will  become  more 
soluble.  This  is  the  commoner  case,  as  is  shown  by  the  way  in 
which  most  solubility  curves  (Fig.  79,  p.  191)  ascend  with  rising 
temperature.  Conversely,  a  substance  which  gives  out  heat  when 
dissolving  in  a  solution  already  almost  saturated  with  the  compound 
is  less  soluble  with  rising  temperature.  For  example,  anhydrous 
sodium  sulphate  gives  out  heat  in  dissolving,  and  so  its  solubility 
diminishes  (Fig.  80,  p.  193)  with  rising  temperature. 

Again,  the  vaporization  of  a  liquid  absorbs  heat,  and  so  an  in- 
crease in  temperature  will  increase  the  pressure,  and  therefore  the 
concentration  of  its  vapor  (p.  146). 

Le  Chatelier's  Law.  —  The  above  mentioned  law  is  really  a 
particular  case  of  a  more  general  one.  If  some  stress  (e.g.,  by  change 
of  temperature,  pressure,  or  concentration)  is  brought  to  bear  on  a 
system  in  equilibrium,  the  equilibrium  is  displaced  in  the  direction 
which  tends  to  undo  the  effect  of  the  stress.  Thus,  raising  the 
temperature  furthers  the  change  which  absorbs  heat  —  and  there- 
fore would  tend  to  lower  the  temperature.  Increasing  the  concentra- 
tion of  the  molecules  pushes  the  action  in  the  direction  which  uses  up 
these  very  molecules  (p.  291).  Again  pressure  causes  ice  to  melt, 
because  the  water  which  is  formed  occupies  a  smaller  volume,  and 
this  change  tends  to  relieve  the  pressure.  But  pressure  will  not 
cause  most  substances  to  melt,  because  usually  the  liquid  form  oc- 
cupies a  greater  volume  and  its  production  would  tend  to  increase 
pressure. 

Summary.  —  In  this  chapter  three  questions  are  answered: 

1.  Why  do  some  chemical  actions  cease,  while  still  incomplete? 
Answer:  They  are  reversible. 

2.  What  explains  the  position  of  the  equilibrium  point?    An- 
swers:  (a)  Equal  effects  of  opposed  molecular  actions;  (6)  Equality 
in  speed  of  opposed  reactions. 

3.  What   will   displace   the   equilibrium   point?    Answer:     (a) 
Change  in  concentration  of  one  (or  more)  of  the  substances;    (b) 
Change  in  the  temperature. 


308  INORGANIC  CHEMISTRY 

Exercises.  —  1.  Why  is  the  formation  of  the  following  sub- 
stances complete:  (a)  silver  chloride  (p.  20),  (6)  hydrogen  chloride, 
and  (c)  water  by  union  of  the  elements? 

2.  How  could  the  interaction  of  chlorine  and  water  (p.  223)  be 
brought  to  completion? 

3.  Explain   why   the   decomposition   of   potassium   chlorate   is 
complete. 

4.  In  view  of  the  statement  on  p.  80,  explain  why  mercuric  oxide 
is  completely  decomposed  by  heating.     Point  out  the  resemblance 
between  this  experiment  and  the  imaginary  one  illustrated  in  Fig.  97 
(p.  302). 

5.  Why  can  magnetic  oxide  of  iron  be  reduced  completely  by  a 
stream  of  hydrogen  (p.  127),  and  iron  oxidized  completely  by  a  current 
of  steam  (p.  116)? 

6.  With  the  phosphorus  pentachloride  system,  say  at  250°,  what 
effect  would  suddenly  enlarging  the  space  containing  a  given  amount 
of  the  vapor  produce?     What  would  be  the  effect  of  diminishing  the 
space?     What  would  be  the  effect  of  introducing  additional  chlorine 
into  the  same  space  (p.  299)? 

7.  By  what  practical  means  could  the  degree  of  dissociation  of 
sulphur  vapor  (Ss)  be  reduced,  without  changing  the  temperature 
(p.  261)? 

8.  What  inference  should  you  draw  from  the  fact  that:   (a)  the 
solubilities  of  potassium  nitrate,  sodium  chloride,  and  Glauber's  salt 
(p.  191)  increase  with  rise  in  temperature  (p.  305);   (6)  that  those  of 
calcium  hydroxide  (p.  190)  and  triethylamine  decrease  with  rise  in 
temperature? 

9.  Is  the  heat  of  solution  of  lead  nitrate  (p.  191)  positive  or  nega- 
tive? 

10.  Carry  out  the  calculation  of  S  for  4  and  8  hours  (p.  297). 

11.  What  is  the  molecular  concentration  of  the  oxygen  in  the  air 
(pp.  9,  294),  of  the  nitrogen  in  the  air,  of  the  aqueous  vapor  above 
water  at  10°  and  at  20°  (p.  146),  of  a  solution  containing  one  formula- 
weight  of  sodium  chloride  in  10  liters,  of  a  solution  containing  65  g. 
of  hydrogen  iodide  in  250  c.c.? 

12.  What  are  the  partial  pressures  of  the  three  components  of 
phosphorus  pentachloride  vapor  at  250°  and  760  mm.   (p.  260)? 
What  are  their  molecular  concentrations? 

13.  Using  the  model  on  p.  298,  study  the  dissociation  of  KI3 
(p.  276),  of  iodine  vapor  (p.  276),  and  of  hydrogen  iodide  (p.  306), 
and  the  formation  of  ferric  thiocyanate  (p.  292).     Show  in  each  case 


CHEMICAL  EQUILIBRIUM  309 

the  effect  on  the  system  of  increase  in  volume  without  change  in 
the  amount  of  material  (p.  299). 

14.  What  actions  in    Chap.  XIV  are  complete  for  the  same 
reason  that  the  action  of  sulphuric  acid  on  salt  (pp.  207-209)  is  so? 

15.  At  a  given  temperature,  would  increasing  the  pressure  in  a 
mixture  of  hydrogen  and  bromine  vapor  render  the  union  more  or  less 
complete?     Is  the  action  more  complete  at  a  high  or  at  a  low  tem- 
perature? 

16.  How  could  a  hydrate  be  completely  dehydrated? 

SUMMARY  OF  PRINCIPLES 

The  summary  of  some  of  the  chief  principles  of  the  science  (p.  229)  may  now 
receive  several  important  additions.  For  the  sake  of  completeness,  reference  tc 
the  periodic  system  is  made  in  No.  21,  to  isomers  in  No.  22,  and  to  the  phase  rule 
in  No.  23,  although  these  subjects  have  not  yet  been  taken  up. 

15.  That  weight  of  each  substance  which  in  the  gaseous  condition  occupied  the 
same  volume  as  32  grams  of  oxygen,  temperature  and  pressure  being  alike  for 
both  (namely,  22.4  liters  at  0°  and  760  mm.),  is  taken  as  the  chemical  unit  of 
weight  for  the  substance,  and  is  known  as  its  molar  weight  (p.  236). 

16.  That  weight  of  each  element  which  is  the  greatest  common  measure  of 
the  quantities  of  the  element  found  in  the  molar  weights  of  its  compounds  is  taken 
as  the  chemical  unit  of  weight  for  the  element,  and  is  known  as  its  atomic  weight. 
This  weight  has  the  property  described  in  6  (p.  230). 

The  composition  of  each  substance  is  expressed  in  terms  of  the  atomic  weights 
as  units,  and  the  sum  of  the  atomic  weights  is  multiplied  by  an  integer,  when 
necessary,  so  as  to  equal  the  molar  weight  (p.  249). 

17.  The  number  of  equivalent  weights  of  hydrogen  which  combine  with,  or  are 
replaced  by  the  atomic  weight  of  an  element  is  called  the  valence  of  the  element 
(p.  132). 

18.  The  speed  of  every  interaction  is  a  function  of  the  first,  or  some  higher 
power  of  the  molar  concentration  of  each  interacting  substance  (p.  294). 

19.  Substances  undergoing,  at  a  fixed  temperature,  an  interaction  which  is 
reversible,  reach  a  condition  of  equilibrium.     The  final  proportions  of  the  mate- 
rials are  such  that  the  speeds  (see  18)  of  the  opposed  actions  are  equal  (p.  290). 

20.  Van't  Hoff's  law  and  Le  Chatelier's  law  (pp.  305,  307). 

21.  Each  element  has  its  own  set  of  chemical  relations  (pp.  226,  284):  e.g.,  it 
can  exist  in  combination  with  certain  other  elements;   it  has  a  certain  valence, 
and  may  have  more  than  one  valence;   it  confers  certain  properties  on  its  com- 
pounds as  a  class;    it  is  metallic  or  non-metallic  (pp.  150,  284);    it  resembles 
certain  other  elements  in  several  of  these  respects  (e.g.,  the  halogens),  and  differs 
from  others,  in  a  way  more  or  less  definitely  described  by  its  place  in  the  periodic 
system  (q.v.). 

In  complex  cases,  the  inter-relations  of  the  elementary  units  in  a  compound, 


310  INORGANIC  CHEMISTRY 

and  the  relations  of  the  compound  to  other  compounds  (see  No.  22),  are  repre- 
sented graphically  by  formulae  based  upon  an  hypothesis  of  molecular  structure 
(pp.  263,  322). 

22.  Identical  combinations  of  matter  may  constitute  more  than  one  com- 
pound substance  (isomers,  see  Urea).     These  may  have  equal  molar  weights  (op- 
tical and  structural  isomers),  or  they  may  have  different  molar  weights  (polymers, 
pp.  250,  282). 

23.  In  a  system  in  equilibrium  the  number  of  components  plus  two  equals  the 
number  of  phases  plus  the  number  of  degrees  of  freedom  (Phase  rule,  q.v.). 


CHAPTER  XVI 
OZONE  AND  HYDROGEN  PEROXIDE 

A  FRESH,  penetrating  odor,  resembling  that  of  very  dilute  chlorine, 
was  noticed  by  van  Marum  (1785)  near  an  electrical  machine  in 
operation.  Schonbein  (1840)  showed  that  the  odor  was  that  of  a 
distinct  substance,  which  he  named  ozone  (Gk.  o£«v,  to  smell),  and 
he  discovered  a  number  of  ways  of  obtaining  it.  It  is  questionable 
whether  there  is  any  ozone  in  the  air,  excepting  -temporarily  in  the 
immediate  neighborhood  of  a  natural  or  artificial  discharge  of  elec- 
tricity. 

Formation  of  Ozone.  —  The  most  significant  way  in  which 
ozone  is  formed  is  by  heating  oxygen.  The  proportion,  at  equi- 
librium, increases  as  the  temperature  rises.  This  shows  that  it  is 
formed  with  absorption  of  heat  (van't  HcfFs  law,  p.  305). 

3O2  +  61,400  cal.<=±208. 

The  percentages  of  ozone  formed  are:  at  1296°,  0.1  per  cent;  at 
2048°,  1.52  per  cent;  at  4500°,  16.5  per  cent.  If  the  mixture  is 
allowed  to  cool  slowly,  the  proportion  diminishes  as  the  tempera- 
ture falls,  by  reversal  of  the  above  reaction  until,  at  300°  or  below, 
the  amount  is  practically  zero.  Rapid  cooling,  however,  to  room 
temperature,  at  which  this  reaction  is  very  slow,  will  preserve  most 
of  it.  A  convenient  way  of  demonstrating  its  formation  by  heating 
is  to  immerse  a  platinum  wire,  heated  white  hot  by  an  electric  cur- 
rent, or  a  small  jet  of  burning  hydrogen,  under  the  surface  of  some 
liquid  air.  The  ozone  is  formed  close  to  the  hot  wire  or  flame,  and 
is  instantly  cooled  as  it  leaves  that  region  by  contact  with  the  liquid 
air  (—190°),  and  so  1.5-2  per  cent  of  it  is  found  in  the  gases  evapo- 
rated by  the  heat.  A  recognizable  trace  of  ozone  is  even  formed 
when  a  small  jet  of  oxygen  is  blown  through  the  tip  of  a  Bunsen 
flame. 

Ozone  is  found  in  the  oxygen  generated  by  electrolysis  of  dilute 
sulphuric  acid  (p.  120).  It  arises  during  the  slow  oxidation  of  phos- 
phorus by  the  air,  resulting,  probably,  from  the  decomposition  of 

311 


312  INORGANIC  CHEMISTRY 

unstable,  highly  oxidized  bodies  which  are  formed  during  the  action. 
Oxygen  containing  as  much  as  15  per  cent  of  it  is  produced  by  the 
interaction  of  fluorine  and  water  (p.  281). 

Preparation  of  Ozone.  —  The  most  satisfactory  way  of  prepar- 
ing ozone  O3  is  to  furnish  the  necessary  energy  by  allowing  electric 
waves  to  pass  through  oxygen.  The  apparatus  (Fig.  98)  consists 
of  two  co-axial  glass  tubes,  between  which  the  oxygen  flows.  The 


^4 


v 


FIG.  98. 

waves  are  generated  by  connecting  an  outer  layer  of  tinfoil  on  the 
outer  tube,  and  an  inner  layer  of  tinfoil  in  the  inner  tube  with  the 
poles  of  an  induction  coil.  With  dry,  cold  oxygen,  about  7.5  per 
cent  of  the  gas  is  turned  into  ozone.  Etching  the  surface  of  the 
glass  next  to  the  oxygen  with  hydrogen  fluoride  improves  the  yield. 

Physical  Properties  of  Ozone.  —  Ozone  is  a  gas  of  blue  color. 
It  boils  at  — 119°,  so  that  when  a  mixture  of  oxygen  and  ozone  is  led 
through  a  U-tube  immersed  in  liquid  oxygen  (  —  182.5°),  the  ozone 
is  liquefied.  The  deep-blue  fluid  contains  only  about  14  per  cent  of 
oxygen,  and  this  may  be  removed  by  evaporation.  When  this  liquid 
is  distilled,  the  last  portion  contains  oxozone  04  (C.  Harries),  which 
constitutes  about  11  per  cent  of  the  ozone  made  by  the  use  of  electric 
waves. 

Ozone  is  much  more  soluble  in  water  than  is  oxygen.  At  12°, 
100  volumes  of  water  would  dissolve  50  volumes  of  the  gas  at  one 
atmosphere  pressure.  Its  solubility,  when  mixed  with  oxygen,  is 
in  proportion  to  its  partial  pressure  (p.  189). 

Chemical  Properties  of  Ozone.  —  Ozone  can  be  kept  unde- 
composed  only  when  mixed  with  much  oxygen.  Hence  its  density 
and  molar -weight  cannot  be  ascertained  save  by  indirect  means. 
The  weight  of  a  liter  of  the  mixture  at  0°  and  760  mm.  having  been 
measured,  the  ozone  may  be  removed  by  absorption  in  turpentine 
and  the  volume  of  it  present  in  the  gaseous  mixture  be  thus  ascer- 


OZONE  AND  HYDROGEN  PEROXIDE  313 

tained.  For  example,  if  the  weight  of  1  1.  was  1.468  g.  and  50  c.c. 
were  absorbed  by  turpentine,  there  were  950  c.c.  of  oxygen.  The 
weight  of  this  oxygen  is  1000  :  950  ::  1.429  :  x,  from  which  x  = 
1.361  g.  The  rest  of  the  weight,  1.468  -  1.361  or  0.107  g.,  was 
that  of  50  c.c.  of  ozone.  The  weight  of  1  1.  of  ozone  at  0°  and  760 
mm.  is  therefore  2.140  g.  The  molecular  weight  (weight  of  22.4  1.) 
is  thus  47.9  g.,  or  nearly  48  g.  The  formula  of  ozone  is  therefore  03. 
When  produced  in  cold  oxygen,  by  energy  from  electric  waves, 
it  decomposes  slowly.  But  this  change,  like  all  others,  is  hastened 
by  raising  the  temperature.  Equilibrium,  with  almost  no  ozone,  is 
reached  instantly  at  250-300°.  Liquid  ozone  sometimes  decom- 
poses explosively.  As  the 'equation  shows: 

3O2<=±2O3 

three  volumes  of  oxygen  give  two  volumes  of  ozone.  That  this 
equation,  showing  that  three  molecules  of  oxygen  give  two  mole- 
cules of  ozone,  is  correct,  may  be  demonstrated  by  measuring  the 
diminution  in  volume  which  accompanies  the  action.  If  a  shrink- 
age of  5  c.c.  is  observed  in  forming  the  ozone,  it  is  found  that  10  c.c. 
more  are  then  absorbed  by  turpentine.  Thus  the  ozone  occupied 
10  c.c.,  and  the  total  oxygen  from  which  it  was  made  was  therefore 
15  c.c.  Hence  three  volumes  of  oxygen  give  two  of  ozone. 

Ozone  is  a  much  more  active  oxidizing  agent  than  is  oxygen. 
Mercury  and  silver,  which  are  not  affected  by  the  latter,  are  con- 
verted into  oxides  by  the  former.  Silver  gives  the  peroxide,  Ag2O2: 

2Ag  +  2O3  ->  Ag202  +  202, 

and  this  action  is  used  as  a  test  for  ozone.  Paper  dipped  in  starch 
emulsion  containing  a  little  potassium  iodide  is  also  used  as  a  test: 

O3  +  2KI  +  H2O  ->  02  +  2KOH  +  I2. 

The  iodine  gives  a  deep-blue  color  to  the  starch  (cf.  p.  276).  This 
test,  however,  will  not  distinguish  ozone  from  chlorine  or  hydrogen 
peroxide,  and  may,  therefore,  be  used  only  in  the  absence  of  these 
substances.  The  last  substance  is  always  present  in  the  air  and, 
since  air  usually  shows  the  above  action,  is  probably  responsible  for 
the  belief  that  air  contains  ozone.  The  action  on  silver  has  never 
been  obtained  with  air. 

Ozone  also  removes  the  color  from  many  of  the  vegetable  color- 
ing matters  and  artificial  dyes.  It  should  be  understood  that  the 
great  majority  of  the  complex  compounds  of  carbon  are  colorless. 


314  INORGANIC  CHEMISTRY 

Even  a  slight  chemical  change,  affecting  only  one  or  two  of  the 
atoms  in  a  complex  molecule,  is  thus  almost  sure  to  give  a  color- 
less or  much  less  strongly  colored  material.  Indigo,  Ci6Hi0N2O2, 
which  has  a  deep-blue  color,  is  an  example  of  a  vegetable  dye  that 
is  also  made  artificially.  When  ozonized  air  is  bubbled  through  a 
dilute  solution  of  this  dye  (as  indigo-carmine),  the  indigo  is  oxidized 
to  isatin  C8H5NO2,  and  the  color  disappears  (see  below). 

Oxygen  and  ozone  are  different  substances  (p.  4),  that  is,  have 
different  properties.  The  difference  in  density,  interpreted  in  terms 
of  the  molecular  hypothesis,  gives  us  the  statement  of  the  nature  of 
the  difference  which  is  embodied  in  the  formulae  02  and  03.  The 
difference  in  activity,  interpreted  in  terms  of  the  conception  of  en- 
ergy, gives  us  the  other  method  of  stating  the  nature  of  the  difference. 
The  recent  preference  for  the  second  method  is  well  illustrated  by 
this  case.  The  first  method  uses  a  mere  physical  property,  the 
second  a  fact  which  is  intimately  connected  with  the  whole  chemical 
behavior  of  the  substance,  a  matter  of  much  greater  interest  to  the 
chemist. 

Ozone  may  be  distinguished  from  chlorine,  nitrogen  peroxide,  and  other  oxi- 
dizing agents,  with  the  exception  of  hydrogen  peroxide,  by  using  pink  litmus  paper 
instead  of  plain  paper  to  carry  the  potassium  iodide  solution  in  the  above  test. 
The  potassium  hydroxide  set  free  by  ozone  turns  the  paper  blue.  Chlorine,  for 
example,  gives  an  entirely  different  action:  CU  +  2KI  — >•  2KC1  +  I2. 

Ozone  is  used  commercially  in  bleaching  oils,  waxes,  ivory,  flour, 
and  starch.  It  is  employed  also  for  sterilizing  drinking  water  in 
Petrograd,  Lille,  and  other  cities.  For  this  purpose,  however, 
bleaching  powder  is  less  expensive.  Ozone  is  used  also,  in  ventila- 
tion, to  destroy  (or  obscure)  the  odors  in  the  animal  houses  of  zoologi- 
cal gardens,  and  to  kill  the  bacteria  and  spores  carried  by  the  dust  in 
the  air.  A  rather  high  concentration  is  required  for  the  last-named 
purpose  however. 

Oxidizing  Agents ,   and  Explanation   of  their  Activity.  — 

When  ozone  turns  into  oxygen  much  heat  is  liberated ~  (equation, 
p.  311).  Ozone  possesses,  therefore,  much  more  internal  energy 
than  does  oxygen.  On  this  account  it  brings  to  the  task  of  oxidiz- 
ing any  substance  more  energy  than  does  oxygen  itself,  and  is  there- 
fore more  efficient.  Thus,  free  oxygen  does  not  interact  in  the  cold 
with  indigo,  or  with  silver  or  potassium  iodide  (see  above),  while 
ozone  oxidizes  them  rapidly. 


OZONE  AND  HYDROGEN  PEROXIDE  315 

The  heats  of  reaction  show  the  difference  very  clearly.  In 
equation  (2),  below,  1800  cal.  is  the  amount  of  heat  which  would  be 
liberated  if  indigo  could  be  oxidized  to  isatin  by  oxygen  gas.  When 
ozone  is  used,  we  obtain,  in  addition,  the  heat  of  decomposition  of 
this  substance  (equation  1),  so  that  the  total  heat  liberated  (equa- 
tion 3),  63,200  cal.,  is  35  times  as  great  as  in  equation  (2)  where  free 
oxygen  is  the  oxidizing  agent: 

203  =  202(+2O)        +      61,400  cal.  (1) 

Ci6H10N2O2  +  (2O)  =  2C8H5NO2          -f        1,800  cal.  (2) 

Ci6H10N202  +  203  =  2C8H5NO2  +  2O2  +  63,200  cal.  (3) 

By  similar  reasoning  we  explain  the  superiority  of  potassium  per- 
manganate over  free  oxygen  for  oxidizing  hydrochloric  acid  (p. 
218).  Substances  which  are  more  active  oxidizers  than  is  free 
oxygen  belong  to  the  class  of  oxidizing  agents. 

It  should  be  noted  that  when  ozone  acts  as  an  oxidizing  agent, 
usually  only  one  of  the  atoms  of  oxygen  in  each  molecule  plays  this 
part,  and  oxygen  gas  is  formed.  This  is  illustrated  in  all  the  three 
examples  cited  in  the  preceding  section. 

Allotropic  Modifications.  —  We  have  seen  that  a  substance 
may  exist  in  more  than  three  regular  states,  solid,  liquid,  and  gaseous. 
When  a  simple  substance  shows  more  than  one  form,  in  the  same 
state,  like  oxygen  and  ozone,  we  call  them  allotropic  modifications. 

HYDROGEN  PEROXIDE  H202 

Preparation  of  Hydrogen  Peroxide.  —  When  sodium  peroxide 
(q.v.)  is  added,  a  little  at  a  time,  to  a  dilute  acid,  hydrogen  peroxide 
is  set  free  and  remains  dissolved  in  the  water: 

Na*O2  +  2HC1  fc?  2NaCl  +  H202. 

When  hydrated  barium  peroxide  Ba02,8H2O  is  shaken  with  cold, 
dilute  sulphuric  acid  a  similar  action  takes  place: 

Ba02  +  H2SO4  *=>  BaS04  1  +  H202. 


The  substance  was  discovered  by  Thenard  (1818)  by  the  use  of  this 
reaction.  The  excess  of  sulphuric  acid  may  be  removed  by  adding 
barium  hydroxide  solution  cautiously  until  no  further  precipitation  of 
barium  sulphate  occurs:  Ba(OH)2  +  H2SO4  *=»BaSO4  1  +  2H2O.  Hy- 
drochloric acid  or  phosphoric  acid  may  be  used  instead  of  sulphuric 


316  INORGANIC  CHEMISTRY 

acid.  The  second  is  largely  employed  in  the  commercial  manu- 
facture of  hydrogen  peroxide.  In  each  case,  great  care  has  to  be 
taken  to  precipitate  the  other  products  and  all  impurities  from  the 
solution.  When  hydrochloric  acid  is  used,  for  example,  the  barium 
chloride  produced  by  the  action  is  removed  by  adding  silver  sulphate  : 


BaCl2  +  Ag2SO4^  BaSO4  1  +  2AgCl  |  . 

An  aqueous  solution  is  also  obtained  by  passing  carbon  dioxide 
through  barium  peroxide  suspended  in  water: 

C02  +  H2O  <=±  H2C03  +  Ba02  *=?  BaC03  1  +  H2O2. 

Pure  hydrogen  peroxide  is  isolated  from  any  of  these  solutions  by 
distillation  under  reduced  pressure.  To  secure  the  low  pressure,  the 
ordinary  distilling  apparatus  (Fig.  20,  p.  43)  is  made  completely 
air-tight,  and  is  connected  by  a  branch  tube  with  a  water-pump. 
Hydrogen  peroxide  is  much  less  volatile  than  water,  but  decomposes 
into  water  and  oxygen  violently  at  100°.  Hence  the  lower  pressure 
is  required  to  make  possible  its  volatilization  at  a  temperature  below 
this  point.  At  26  mm.  pressure,  the  water  begins  to  pass  off  first 
(at  about  27°).  The  last  portion  of  the  liquid  boils  at  69°  and  is 
almost  all  hydrogen  peroxide. 

By  evaporating  the  commercial  (3  per  cent)  solution  at  70°,  a 
liquid  containing  45  per  cent  of  hydrogen  peroxide  may  be  made 
without  much  loss  of  the  material  by  volatilization. 

Hydrogen  peroxide  was  formerly  separated  from  the  other  sub- 
stances produced  in  the  reaction  for  its  preparation,  and  from  a 
large  part  of  the  water,  by  repeatedly  shaking  the  mixture  with 
ether  (cf.  p.  180).  The  relative  solubility  in  water  and  ether  is 
1  :  0.0596,  however,  so  that  much  ether  is  needed.  The  ethereal 
layer,  which  rises  to  the  top,  when  evaporated,  leaves  a  strong 
aqueous  solution  of  the  compound  behind.  Explosive  substances 
are  often  formed  by  interaction  with  the  ether  (perhaps  ethyl 
peroxide  (C2H5)202),  however,  and  so  this  method  is  no  longer 
employed. 

The  Interaction  of  Barium  Peroxide  and  Sulphuric  Acid. 

—  It  is  worth  noting  that,  although  common  barium  peroxide  is  not 
less  soluble  in  water  than  is  the  hydrated  form  used  above,  it  dissolves 
much  more  slowly.  The  fact  that  it  is  made  by  heating  barium 
oxide  in  oxygen  and  is  composed  of  compact  particles  is  perhaps 
accountable  for  this. 


OZONE  AND  HYDROGEN  PEROXIDE  317 

Every  action  upon  a  little-soluble,  or  slowly  dissolving  body,  like 
the  barium  peroxide  in  the  above  actions,  is  rather  complex.  It  is 
only  the  dissolved  part  of  the  substance  that  interacts.  There  is 
thus  a  physical  equilibrium  between  the  undissolved  and  the  dissolved 
bodies,  BaO2  (solid)  ^  BaO2  (dslvd),  the  displacement  of  which 
furnishes  the  material  for  the  chemical  action.  The  latter  has  there- 
fore to  follow  the  pace  set  by  the  former.  When  barium  sulphate 
is  precipitated,  another  physical  equilibrium  follows  the  chemical 
change :  BaS04  (dslvd)  ^±  BaSO4  (solid) .  When  relatively  insol- 
uble bodies  are  used  or  produced,  there  is  thus  a  chain  of  equilibria 
each  depending  on  the  others: 
BaO2  (solid)  *=5  BaQ2  (dslvd)  +H2SO4£^  H2O2  +  BaSO4  (dslvd)  ±=?  BaSO4  (solid). 

If  the  barium  sulphate  ceased  to  be  precipitated,  its  interaction  in 
solution  with  the  hydrogen  peroxide  would  drive  the  central  action 
backwards,  and  barium  peroxide  would  be  precipitated  instead. 
The  success  of  the  process  thus  depends  on  the  fact  that  barium  sul- 
phate is  less  soluble  than  is  barium  peroxide. 

When  carbon  dioxide  is  used  (see  above),  a  similar  chain  of  equi- 
libria exists,  and  in  that  case  it  is  the  barium  carbonate  that  is  the 
less  soluble  substance. 

Other  Modes  of  Formation.  —  Hydrogen  peroxide  is  found  in 
minute  amounts  in  rain  and  snow.  It  is  formed  by  the  direct  union 
of  hydrogen  and  oxygen.  When  a  hydrogen  flame  is  allowed  to 
play  upon  ice,  appreciable  amounts  of  the  peroxide  are  saved  from 
being  decomposed,  as  they  ordinarily  would  be  by  the  heat  of  the 
action,  and  are  found  in  the  water.  It  is  formed  also  to  the  extent 
of  37  per  cent  when  electric  waves  pass  through  a  mixture  of  hydro- 
gen and  oxygen  cooled  by  liquid  air  (—190°).  It  is  produced  when 
oxygen  is  passed,  close  to  the  negative  electrode,  through  the  liquid  in 
an  electrolytic  cell  containing  dilute  sulphuric  acid.  The  gas  is  re- 
duced by  the  hydrogen  being  liberated  on  the  platinum  plate. 

Traces  of  hydrogen  peroxide  are  formed  when  zinc,  copper,  lead, 
and  other  metals  are  shaken  with  air  and  very  dilute  sulphuric 
acid  (1  :  55  Aq).  It  has  been  suggested  that  the  water  loses  its 
oxygen  to  the  zinc  and  gives  its  hydrogen  to  the  oxygen: 


O2  +  H2  O  +  Zn 


H2O2  4-  ZnO  +  H2SO4  ->  ZnSO4  +  H2O, 


the  action  being  assisted  by  the  tendency  of  the  zinc  oxide  to  act 
with  the  sulphuric  acid  to  give  zinc  sulphate. 


318  INORGANIC  CHEMISTRY 

Physical  Properties.  —  Hydrogen  peroxide  is  a  syrupy  liquid 
(density  1.46).  It  blisters  the  skin,  and,  when  diluted,  has  a  dis- 
agreeable metallic  taste.  It  has  been  frozen  (m.-p.  —0.8°),  and 
boils  at  26  mm.  pressure  at  69°  and  at  46  mm.  at  80°. 

Chemical  Properties.  —  Hydrogen  peroxide  (100  per  cent)  is 
very  unstable,  and  decomposes  slowly  even  at  —20°.  The  dilute 
aqueous  solution,  when  free  from  impurities,  keeps  fairly  well.  The 
presence  of  a  trace  of  free  acid  increases  its  stability.  Free  alkalies 
and  most  salts  assist  the  decomposition;  hence  the  necessity  for 
purifying  the  commercial  solution.  Since  free  acids  must  not  be 
used  to  stabilize  solutions  for  medical  use,  a  trace  of  some  organic 
compound  which  has  the  same  effect  (several  such  are  known)  is 
added.  Addition  of  powdered  metals,  of  manganese  dioxide,  and 
of  charcoal  causes  effervescence  even  in  dilute  solutions,  and  oxygen 
escapes: 

2H2O2-»2H2O  +  02. 

The  3  per  cent,  commercial  solution  yields  in  this  way  ten  times  its 
own  volume  of  oxygen,  and  so  is  often  labelled  "10  vol.  solution." 
The  more  concentrated  solutions  (38  per  cent)  remain  quiescent  in  a 
dish  of  polished  platinum  even  at  60°,  but  the  making  of  a  slight 
scratch  on  the  bottom,  beneath  the  surface  of  the  liquid/  causes  pro- 
fuse liberation  of  oxygen  along  the  sharp  edge  thus  produced.  The 
action  of  these  contact  agents  is  therefore  probably  mechanical. 

Since  the  substance  cannot  be  vaporized,  even  at  low  pressure, 
without  some  decomposition,  its  molar  weight  has  been  determined 
by  the  freezing-point  method  (p.  199).  The  freezing-point  of  a  3.3 
per  cent  solution  in  water  is  2.03°  below  that  of  the  water  itself. 
Hence,  in  1000  g.  of  water,  3.3  g.  would  have  given  a  depression  of 
2.03  X  96.7  -f-  1000,  or  0.196°.  Therefore  a  depression  of  1.86°,  the 
average  depression  produced  by  one  mole  of  a  substance  in  1000  c.c. 
of  water  (see  p.  335),  would  have  been  caused  by  3.3  X  1.86  -f-  0.196, 
or  31.3  g.,  which  is  the  required  molar  weight.  Now  the  formula 
HO  corresponds  to  a  molar  weight  of  17  and  H2O2  to  one  of  34.  It 
is  evident,  therefore,  that  the  latter  is  the  correct  formula. 

Hydrogen  peroxide,  in  solution  in  water,  is  a  feeble  acid.  The 
normal  molar  weight  and  very  small  electrical  conductivity  (see 
Chap.  XVIII)  show  that  only  a  very  small  proportion  of  it  can  be 
ionized.  As  an  acid  it  enters  into  double  decomposition  readily  and 
the  peroxides  are  salts  with  the  negative  radical  O2"  (peroxidates). 


OZONE  AND  HYDROGEN  PEROXIDE  319 

Thus,  when  it  is  added  to  solutions  of  barium  and  strontium  hy- 
droxides, the  hydrated  peroxides  appear  as  crystalline  precipitates: 

Sr(OH)2  +  H2O2  <±  2H20  +  Sr02. 

The  precipitation  involves  another  equilibrium:  SrO2  +  8H20  <=± 
SrO2,8H20  (solid). 

The  formation  of  a  beautiful  blue  substance  by  the  action  of 
hydrogen  peroxide  upon  dichromic  acid  is  used  as  a  test.  The  test 
is  carried  out  by  adding  a  drop  of  potassium  dichromate  to  an  acidu- 
lated solution  of  the  peroxide.  The  acid  interacts  with  the  dichro- 
mate, giving  free  dichromic  acid:  - 

H2S04  +  K2Cr2O7  <=±  H2Cr207  +  K2S04. 

The  blue  substance,  which  is  very  unstable  and  quickly  decom- 
poses, is  a  perchromic  acid.  A  blue,  crystalline  perchromic  acid 
(HO)4Cr(OOH)3,  which  decomposes  above  —30°,  has  been  prepared. 
The  blue  substance  has  the  property,  unusual  in  inorganic  compounds, 
of  dissolving  much  more  readily  in  ether  than  in  water.  It  is  also 
much  less  unstable  when  removed  from  the  foreign  materials  in  the 
aqueous  solution.  Hence  the  test  is  rendered  more  delicate  by  ex- 
tracting the  solution  with  a  small  amount  of  ether.  In  the  ethereal 
layer  the  color  of  the  compound  is  more  permanent,  as  well  as  more 
distinctly  visible  on  account  of  the  greater  concentration. 

Hydrogen  peroxide  is  a  much  more  active  oxidizing  agent  than 
is  free  oxygen.  This  would  be  expected  from  the  fact  that  it  con- 
tains so  much  more  energy  than  the  water  and  oxygen  into  which 
it  decomposes  (p.  318):  H202  ->  H2O  +  O"+  23,100  cal.  Thus,  it 
liberates  iodine  from  hydrogen  iodide,  an  action  which,  in  presence 
of  starch  emulsion  (cf.  p.  276),  is  used  as  a  test  for  its  presence: 


It  converts  sulphides  into  sulphates.  The  white  lead  (q.v.)  used  in 
paintings  is  changed  by  the  hydrogen  sulphide  in  the  air  of  cities  to 
black  lead  sulphide,  Pb3(OH)2(CO3)2  +  3H2S  -+3PbS  +  4H2O  +  2CO2. 
This  may  be  oxidized  to  white  lead  sulphate  by  means  of  hydrogen 
peroxide: 

PbS  +  4H202  -»  PbS04  +  4H2O, 

and  in  this  way  the  original  tints  of  the  picture  may  be  practically 
restored.  Organic  coloring  matters  are  changed  into  colorless  sub- 
stances by  an  action  similar  to  that  of  ozone  (cf.  p.  314).  Hence 


320  INORGANIC  CHEMISTRY 

hydrogen  peroxide  is  used  for  bleaching  silk,  feathers,  hair,  and 
ivory,  which  would  be  destroyed  by  a  more  violent  agent.  The 
products  of  its  decomposition,  being  water  and  oxygen  only,  are 
harmless,  and,  on  this  account,  it  is  used  in  disinfecting  (destroying 
organisms  in)  infected  sores,  and  as  a  throat  wash. 

Hydrogen  peroxide  exercises  the  functions  of  a  reducing  agent  in 
special  cases,  also.     Thus,  silver  oxide  is  reduced  by  it  to  silver: 

Ag20  +  H202  **  2Ag  +  H20  +  02. 

A  solution  of  potassium  permanganate,  in  which  the  permanganic 
acid  has  been  set  free  by  an  acid,  KMnO4  +  H2S04  *=±  HMnO4  + 
KHS04,  is  rapidly  reduced.  The  permanganic  acid,  with  excess  of 
sulphuric  acid,  tends  to  undergo  the  first  of  the  following  changes, 
provided  a  substance  is  present  ivhich  can  take  possession  of  the  oxygen 
that  would  remain  as  a  balance: 

2HMnO4  +  2H2SO4  ->  2MnS04  +  3H2O  (+  50)  (1) 

(5O)  +  5H202->5H2O  +  502  (2) 


2HMnO4  +  2H2SO4  +  5H2O2  ->  2MnSO4  +  8H20 

In  all  reductions  by  hydrogen  peroxide,  each  molecule  of  the  latter  removes 
but  one  atomic  weight  of  oxygen  from  the  other  substance.  Whether  it  behaves 
thus  because  its  two  hydrogen  units  combine  with  this  oxygen  and  all  its  own 
oxygen  escapes,  or  because  it  furnishes  water  and  one  oxygen  unit  of  the  pair 
required  to  form  the  molecule  of  free  oxygen  (the  substance  reduced  furnishing 
the  other),  has  not  been  determined. 

The  above  action  is  used  in  quantitative  analysis  for  estimating 
the  quantity  of  hydrogen  peroxide  in  a  given  liquid  after  the  liquid 
has  been  acidified.  The  amount  of  a  standard  (p.  277)  solution  of 
the  permanganate  which  is  required  to  decompose  all  the  peroxide  is 
measured  by  means  of  a  burette  (q.v.).  The  permanganate  is  deep 
reddish-purple  in  color,  while  the  products  are  colorless.  Hence, 
after  the  peroxide  is  exhausted,  the  next  drop  of  the  permanganate 
confers  a  distinct,  permanent,  pink  tinge  upon  the  liquid.  The  addi- 
tion of  the  permanganate  solution  is  stopped  so  soon  as  this  condi- 
tion is  reached  and  the  volume  of  it  that  has  been  used  is  read  off. 

The  action  of  hydrogen  peroxide  on  hydrogen  iodide  proceeds  slowly,  so  that 
its  speed  can  be  measured.  Although  the  equation  shows  three  interacting  mole- 
cules (2HI  +  H2O2),  a  constant  (p.  295)  is  obtained  only  by  using  the  formula 
for  a  reaction  involving  two  molecules  (reaction  of  the  second  order).  This  is 
because  the  reaction  takes  place  as  two  consecutive  actions: 

HI  +  H202  =  HIO  +  H20,  (1) 

El  +  HIO  *=±  H20  +  I2,  (2) 


OZONE  AND  HYDROGEN   PEROXIDE  321 

of  which  (1),  in  which  hypo-iodous  acid  is  formed,  alone  takes  much  time. 
The  second  (2)  is  the  reverse  of  the  action  of  iodine  on  water  (cf.  p.  277), 
and  is  very  speedy.  Hence  the  speed  measurement  concerns  only  (1),  which  is 
dimolecular. 

Thermochemistry  of  Hydrogen  Peroxide.  —  The  formation 
of  hydrogen  peroxide  from  the  free  elements  is  accompanied  by  evo- 
lution of  heat : 

H2  +  O2  =  H2O2  Aq  +  45,300  cal. 

Hence  the  substance  is  formed  by  direct  union  (p.  317).  But  its  de- 
composition into  water  and  oxygen  gives  out  a  further  supply  of 
heat: 

H2O2  =  H2O  +  O  +  23,100  cal. 

The  sum  of  these  two  stages,  of  course,  yields  the  same  result  (cf. 
p.  100)  as  the  direct  formation  of  water  (68,400  cal.). 

When  hydrogen  peroxide  is  used,  instead  of  free  oxygen,  for  oxi- 
dizing purposes,  each  such  action  liberates  23,100  calories  of  heat 
more  in  the  former  case  that  it  would  in  the  latter.  Hence  the  activ- 
ity of  the  substance  as  an  oxidizer  (cf.  pp.  314,  319). 

Chemical  Constitution  of  Peroxides.  —  We  have  seen  (p. 
317)  that  when  acids  act  upon  barium  peroxide  BaO2,  hydrogen 
peroxide  is  formed:  Ba02  +  H2S04  ->  BaSO4  +  H202.  But  all 
oxides  containing  two  atoms  of  oxygen  in  each  molecule  do  not 
yield  hydrogen  peroxide.  Thus,  lead  dioxide  Pb02  and  manganese 
dioxide  Mn02,  when  treated  with  sulphuric  acid,  give  the  sulphate 
of  the  metal  and  water  and  oxygen:  2Mn02  -f  2H2SO4  — >  2MnS04+ 
2H2O  +  O2.  We  infer  from  this  that  the  two  sets  of  dioxides  are 
not  alike.  With  barium  peroxide  there  is  a  simple  double  decom- 
position, or  exchange  of  radicals,  and  so  we  hold  that  it  contains  the 
radical  O2  which  is  bivalent  as  a  whole:  BaII(O2)11.  This  harmo- 
nizes with  the  fact  that  no  other  compound  is  known  in  which  barium 
has  even  the  appearance  of  being  quadrivalent.  Manganese  dioxide 
can  give  a  tetrachloride  (p.  219),  however,  as  can  also  lead  dioxide 
(q.v.),  so  we  infer  that  the  manganese  and  lead  are  here  quadrivalent, 
and  that  the  radical  is  O:  MnIV(0)2n  and  PbIV(0)2U.  For  this 
reason,  we  have  recently  begun  to  call  barium  peroxide  a  salt,  and 
to  name  it  barium  peroxidate.  Similarly,  we  have  strontium  perox- 
idate  Sr11^)11  and  sodium  peroxidate  Na2(O2)n.  Hz(02)u  should 
therefore  be  called  peroxidic  acid. 


322  INORGANIC  CHEMISTRY 

There  still  remains  the  question  whether  hydrogen  peroxide  and 
barium  peroxide  are: 

I  II 

H-0  O  H 

I    and  Ba(  |         or  )o  =  O  and  Ba  =  O  =  O. 

H-0  O  H 

In  other  words,  whether  the  negative  radical  of  the  peroxidates  is 
—  0  —  O—  or  =O  =  0.  Some  compounds  in  which  oxygen  is  quad- 
rivalent are  known,  so  that  the  second  alternative  is  worthy  of  con- 
sideration. Now,  we  have  seen  that  hydrogen  peroxide  is  formed 
by  the  reduction  of  dissolved  oxygen  (p.  317).  On  the  whole,  this 
favors  the  symmetrical  formula  I,  rather  than  the  unsymmetrical 
formula  II.  Then,  hydrogen  peroxide  is  never  formed  by  the  oxi- 
dation of  water.  This  favors  I  very  distinctly,  because  it  would  be 
more  difficult  to  insert  the  oxygen  to  give  I  than  to  attach  it  to  the 
oxygen  of  the  water  as  in  II.  Finally,  when  some  ethyl  sulphate 
(62115)2804  is  dissolved  in  15  per  cent  hydrogen  peroxide,  and  sodium 
hydroxide  is  added  a  drop  at  a  time  while  the  mixture  is  shaken,  the 
following  reaction  occurs: 

H202  +  (C2H6)2S04  +  2NaOH  ->  (C2H5)202  +  Na*S04  +  2H2O. 


By  distillation,  the  substance  (C2H5)2O2  is  obtained  (b.-p.  65°  at 
760  mm.).  When  zinc  dust  and  acetic  acid  are  added  to  this  prod- 
uct, hydrogen  is  liberated:  Zn  +  2HCO2CH3  ->  Zn(C02CH3)2  +  H2 
and  the  hydrogen  reduces  the  (C2H5)2O2,  giving  alcohol  C2H50H 
(Baeyer  and  Villiger,  1900).  Now,  the  compound  was  either: 

—  0  C2Hs 

I         or  )  O  =  O. 

—  O  O2ti5 

The  second,  on  reduction,  would  be  expected  to  give  ether  C2H5  — 
0  —  C2H5,  which  is  not  formed,  while  the  former  would  give  2C2H5OH, 
which  is  formed.  Hence  formula  I  is  assigned  to  the  peroxidates. 

Such  a  formula  is  called  a  structural  or  graphic  formula,  because, 
by  means  of  a  construction  or  graph,  it  indicates  the  chemical  reac- 
tions of  the  substance.  It  may  also  indicate  the  way  in  which  the 
parts  of  the  molecule  are  actually  connected,  but  its  primary  pur- 
pose is  to  indicate  chemical  behavior.  Of  course,  the  number  of 
lines  emanating  from  each  symbol  must  correspond  with  the  valence 
of  the  atom  concerned. 


OZONE  AND  HYDROGEN  PEROXIDE         323 

Exercises.  —  1.  What  volume  of  ozone  will  be  taken  up  by  100 
c.c.  of  water  at  12°  from  a  stream  of  oxygen  at  760  mm.  containing 
7.5  per  cent  of  ozone  (p.  312)? 

2.  Formulate  the  action  of  carbon  dioxide  on  barium  dioxide 
(p.  316)  after  the  manner  of  that  of  sulphuric  acid  on  the  same 
substance  (p.  317).     The  dissolving  gas  gives  an  additional  equi- 
librium: C02  (gas)  +  H2O^±H2CO3  (dslvd). 

3.  At  what  temperature  will  a  ten  per  cent  solution  of  hydrogen 
peroxide  freeze  (p.  318)? 

4.  Write  the  thermochemical  equations  for  oxidation  of  indigo 
by  hydrogen  peroxide  (pp.  315,  319). 

5.  How  many  times  its  own  volume  of  oxygen  gas  will  a  4  per 
cent  solution  of  hydrogen  peroxide  give  off  when  treated  with: 
(a)  platinum  powder  (p.  318);    (6)  sulphuric  acid  and  potassium 
permanganate? 

6.  What  per  cent  of  hydrogen  peroxide  does  a  "  12  vol. "  solution 
contain? 


CHAPTER  XVII 
DISSOCIATION  IN  SOLUTION 

THE  employment  of  interacting  substances  in  the  form  of  solu- 
tions is  so  constant  in  chemistry,  and  the  reasons  for  this  are  so 
cogent,  that  we  must  now  resume  the  discussion  of  this  subject  (cf. 
p.  178). 

The  present  chapter  will  be  devoted  to  giving  the  proofs  that 
the  molecules  of  acids,  bases,  and  salts,  in  aqueous  solutions,  are 
actually  dissociated  into  parts  by  the  solvent.  This  will  be  shown 
by  consideration,  successively,  of  certain  peculiarities  in  the  chemical 
behavior,  the  osmotic  pressures,  the  freezing-points,  and  the  boiling- 
points  of  the  solutions  of  these  substances.  We  shall  see  that  these 
parts  coincide  in  composition  with  the  radicals,  and  are  called  ions. 

Some  Characteristic  Properties  of  Acids,  Bases,  and  Salts, 
Shown  in  Aqueous  Solution.  —  Acids  all  contain  hydrogen  (p. 
120).  In  aqueous  solution,  if  soluble,  they  are  sour  in  taste,  they 
turn  blue  litmus  red,  and  their  hydrogen  is  displaced  by  certain 
metals  (p.  118),  and  has  the  properties  of  a  radical.  By  the  last 
statement  is  meant  that  it  very  readily  exchanges  places  with  other 
radicals  in  reversible  double  decompositions  (p.  208).  Amongst  the 
acids  mentioned  have  been:  hydrochloric  acid  HC1,  sulphuric  acid 
H2SO4,  hypochlorous  acid  HC1O,  acetic  acid  HCO2CH3,  and  hydro- 
gen peroxide  H202.  Many  other  bodies,  like  sugar,  kerosene,  and 
alcohol,  also,  contain  hydrogen  but  not  one  of  them  shows  all  of 
these  properties. 

Again,  all  salts  (p.  214)  are  made  up  of  two  radicals,  and  the 
reversible  double  decompositions  into  which  they  enter  with  acids, 
bases,  and  other  salts,  consist  in  exchanges  of  these  radicals.  Other 
substances  may  include  the  same  combinations  of  atoms,  but  in  their 
actions  these  groupings  are  often  disregarded.  Thus,  sodium  chlo- 
ride NaCl  and  silver  nitrate  AgNO3  exchange  radicals  completely 
(p.  20),  and,  in  dilute  solution,  hydrogen  chloride  and  sodium  hydro- 
gen sulphate  do  so  partially  (p.  208).  But  sodium  chloride  and 
nitroglycerine  CsHstNOs^  do  not  interact  at  all.  The  latter  is  not  a 

324 


DISSOCIATION  IN  SOLUTION  325 

salt,  although  it  contains  the  same  proportion  of  nitrogen  to  oxygen 
as  does  any  nitrate. 

All  bases  (p.  149)  contain  hydroxyl  OH  as  a  radical,  combined 
with  some  positive  radical.  Potassium  hydroxide  KOH  is  soluble 
and  active,  zinc  hydroxide  Zn(OH)2  and  many  others,  however,  are 
insoluble.  Bases  all  exchange  radicals  readily  in  double  decomposi- 
tion with  salts  and  acids.  Other  substances,  like  alcohol  C2H5OH, 
may  contain  hydroxyl,  but  do  not  interact  readily  with  salts  like 
NaCl,  and  are  not  bases. 

The  Influence  of  Water  and  Other  Solvents.  —  It  is  chiefly 
in  aqueous  solution  that  these  special  properties  of  acids,  bases,  and 
salts  become  apparent.  Their  behavior  is  often  quite  different  in 
the  absence  of  this  solvent.  If,  for  example,  we  mix  together  dry 
ammonium  carbonate  (NH4)2C03  and  partially  dehydrated,  solid 
cupric  nitrate  Cu(NO3)2,  and  apply  heat,  a  violent  interaction  be- 
gins. An  immense  cloud  of  smoke  and  gas  is  thrown  out  of  the  tube, 
and  the  substance  remaining  is  either  black  or  reddish,  in  parts,  ac- 
cording to  the  proportions  of  the  substances  employed.  The  residue 
contains  black  cupric  oxide  CuO,  and  sometimes  red  cuprous  oxide 
Cu2O.  The  gas  is  tinged  red  by  the  presence  of  nitrogen  tetroxide 
NO2,  while  a  more  careful  examination  would  show  that  it  contained 
carbon  dioxide,  nitrogen,  nitrous  oxide  N20,  water  vapor,  and  per- 
haps still  other  products. 

The  contrast,  when  the  substances  are  dissolved  in  water  before 
being  brought  in  contact  with  one  another,  is  very  great.  A  pale- 
green  precipitate  is  formed  at  once,  and  rapidly  settles  out.  On 
examination,  this  turns  out  to  be  a  carbonate  of  copper  (basic, 
see  under  Copper),  while  evaporation  of  the  solution  furnishes  us 
with  ammonium  nitrate.  There  are  only  two  main  products,  and 
the  essential  part  of  the  action  in  solution  may  be  represented  by 
the  equation: 

(NH4)2C03  +  Cu(N03)2  ->  CuC03 1  +  2NH4NO3. 

In  the  interaction  between  the  dry  substances  the  molecules  are 
completely  disintegrated,  the  whole  change  is  very  complex,  and  it 
takes  a  good  deal  of  time.  In  the  action  in  water,  no  heating  is 
required,  the  substances  are  neatly  broken  apart,  certain  groups  of 
atoms,  which  we  call  radicals,  are  transferred  as  wholes  from  one 
state  of  combination  to  another,  and  the  rearrangement  takes  place 
instantaneously  in  a  machine-like  manner.  Contrasts  like  this  be- 


326  INORGANIC  CHEMISTRY 

tween  the  interactions  of  anhydrous  and  dissolved  bodies  are  very 
common.  Thus,  we  have  had  occasion  (p.  119)  to  mention  the 
difference  between  the  action  of  metals  on  concentrated  and  on 
dilute  sulphuric  acid. 

Many  compounds,  however,  do  not  show  any  change  in  behavior 
when  dissolved  in  water.  Sugar,  for  example,  is,  as  a  rule,  more 
readily  acted  upon  in  the  absence  of  any  solvent.  Then  again,  while 
water  is  not  the  only  solvent  which  has  the  effect  we  have  just  de- 
scribed, the  majority  of  solvents,  if  they  affect  chemical  change  at 
all,  simply  retard  it.  Thus  the  union  of  iodine  and  phosphorus  in 
the  absence  of  a  solvent  takes  place  spontaneously  with  a  violent  evo- 
lution of  heat.  When  the  elements  are  dissolved  in  carbon  bisulphide 
before  being  mixed,  the  action  is  much  milder,  although  the  product 
is  the  same  (phosphorus  tri-iodide) .  The  diminution  in  the  concen- 
tration of  the  ingredients  by  solution  has  simply  decreased  the  speed 
of  the  action  in  the  normal  way  (p.  291).  That  water  and  some 
other  solvents  (e.g.,  alcohol)  have  a  specific  influence  tending  to  in- 
crease the  activity  of  acids,  bases,  and  salts,  shows  that  a  special 
explanation  of  the  phenomenon  must  be  found. 

Summing  up  these  points  we  see  that  the  peculiarity  of  acids, 
bases,  and  salts  in  aqueous  solution  is  that  each  compound  always 
splits  in  the  same  way.  Thus,  cupric  nitrate  always  gives  changes 
involving  Cu  and  N03  and  never  interacts  so  as  to  use  CuN2  and  O3, 
or  Cu02  and  NQ2,  as  the  basis  of  exchange.  Similarly,  dilute  acids 
always  offer  hydrogen  in  exchange,  and  so  nitric  acid  behaves  as  if 
composed  of  H  and  NO3,  and  sulphuric  acid  as  if  composed  of  2H 
and  864,  and  never  as  if  made  up  of  HSO  and  HO3,  or  H2S  and  64. 
The  sour  taste  and  the  effect  upon  litmus  seem  to  be  properties  of 
this  easily  separable  hydrogen,  for  they  are  shown  only  by  acids. 
The  result  is  that  we  can  make  a  list  of  the  units  of  exchange,  such 
as  H,  OH,  NO3,  CO3,  S04,  Cu,  K,  and  Cl,  employed  by  acids,  bases, 
and  salts  in  their  interactions.  The  molecule  of  each  compound 
of  these  classes  contains  at  least  two  of  them.  Even  when  these 
units  contain  more  than  one  atom,  their  coherence  is  as  noticeable 
within  this  class  of  actions,  as  is  the  permanence  of  the  atomic  masses 
themselves  in  all  actions. 

The  question  raised  in  our  minds  is  whether  solution  in  water 
alters  the  character  of  the  molecule  simply  by  producing  a  sort  of 
plane  of  cleavage  in  it  which  creates  a  predisposition  to  a  uniform 
kind  of  chemical  change,  or  whether  it  actually  divides  the  molecules 
into  separate  parts  consisting  of  the  above  units  of  exchange,  and 


DISSOCIATION  IN  SOLUTION  327 

leaves  subsequent  chemical  actions  to  occur  by  cross-combination  of 
these  fragments.  The  fact  that  the  dissolved  substances  can  be 
recovered  by  evaporation  of  the  liquid  does  not  demonstrate  that 
they  have  not  been  changed  temporarily  while  in  solution.  The 
alteration  which  the  water  produces,  whatever  it  be,  will  naturally 
be  reversed  when  the  water  is  removed.  Since  our  question  involves 
nothing  but  the  counting  of  particles,  the  number  of  which  would 
be  much  greater  in  the  event  that  actual  subdivision  of  molecules  is 
the  explanation,  it  can  be  answered  by  a  study  of  the  physical  prop- 
erties of  solutions.  Several  physical  properties  may  be  used,  and 
they  give  concordant  answers  to  the  question.  We  shall  consider 
the  evidence  of  osmotic  pressure,  of  freezing-points,  boiling-points 
(in  this  Chapter),  and  of  conductivity  for  electricity  (Chap.  XVIII). 

OSMOTIC  PRESSURE 

•  In  the  earlier  discussion  of  solution  (p.  184)  the  condition  of  a 
dissolved  substance  was  viewed  as  akin  to  that  of  a  gas.  We  con- 
ceived the  molecules  of  the  dissolved  substance  as  being  distributed 
through  the  space  occupied  by  the  solvent,  as  being  separate  from 
one  another,  and  as  moving  about  independently  of  each  other. 
This  was  because  the  phenomena  of  diffusion  and  osmotic  pressure 
(p.  185)  in  solution  resemble  those  of  diffusion  and  pressure  in  gases. 
The  attempt  to  obtain  by  calculation,  using  this  theory,  the  results 
that  are  observed,  shows  that  this  theory  applies,  as  van't  Hoff 
stated  when  suggesting  it,  only  to  infinitely  dilute  solutions.  It 
gives  a  fairly  satisfactory  explanation  of  the  behavior  of  extremely 
dilute  solutions,  but  not  of  solutions  such  as  are  commonly  em- 
ployed in  chemical  work. 

The  invention  of  a  suitable  hypothesis  for  the  explanation  of 
the  facts  of  osmosis  presents  some  difficulties,  but  the  facts  them- 
selves are  undoubted.  It  will  conduce,  therefore,  to  clearness  if  we 
speak  first  of  some  things  which  may  be  observed  and  are  true,  irre- 
spective of  any  explanation. 

Phenomena  Produced  by  Osmotic  Pressure.  —  In  order  that 
the  osmotic  pressure  (Gk.  OKT/AOS,  impulsion)  may  be  perceived,  a  par- 
tition, which  the  dissolved  molecules  are  unable  to  traverse,  must 
be  interposed  between  the  solution  and  a  contiguous  mass  of  the  pure 
solvent  (Fig.  75,  p.  184).  The  partition  must  be  permeable  by  the 
solvent,  however.  Such  a  partition  is  described  as  semi-permeable. 


328 


INORGANIC  CHEMISTRY 


securely  attached. 


The  general  nature  of  the  phenomena  may  be  seen  by  employing 
a  tube  (Fig.  99),  to  which  a  diffusion  shell,  of  test-tube  form,  is 
It  is  charged  with  sugar  solution,  and  suspended 
in  pure  water.     This  thimble  is  somewhat  per- 
meable by  the  sugar,  but  the  water  traverses  it 
very  easily,  and  so  an  exhibition  of  the  general 
result  of  a  stricter  test  is  obtained  quickly. 

The  water  is  able  to  pass  freely  through  the 
membrane  in  either  direction,  while  the  sugar  is 
not.  As  the  result  of  the  interchange  of  water, 
the  liquid  rises  slowly  but  steadily  in  the  tube. 
The  pure  solvent  always  passes  into  the  solution. 
If,  further,  two  solutions  of  different  concentra- 
tions of  the  same  substance  are  employed,  then, 
invariably,  water  passes  from  the  more  dilute  so- 
lution into  the  more  concentrated  one  through  the 
membrane.  There  is  apparently  a  tendency  for 
the  water  so  to  distribute  itself  that  the  solutions 
may  eventually  become  equal  in  strength.  The 
water  passes  from  a  dilute  solution,  leaving  it 
more  concentrated  than  before,  into  a  more  con- 
centrated solution,  rendering  it  more  dilute. 

These  phenomena  were  first  studied  by  Pfeffer 
(1877),  a  botanist,  who  used  certain  plant  cells 
for  the  purpose.  The  cell  content  included  a 
liquid  containing  various  salts  in  solution,  and  a 
protoplasmic  layer  which  was  not  attached  to  the 
cell  wall.  This  protoplasmic  layer  behaved  like 
a  semi-permeable  membrane.  When  such  cells  were  immersed  in 
a  concentrated  solution  of  any  substance,  the  water  passed  from  the 
interior  of  the  cell  to  the  solution,  and  by  means  of  a  microscope 
a  shrinkage  of  the  protoplasmic  layer  away  from  the  cell  wall  could 
be  observed.  Conversely,  when  such  cells  were  placed  in  pure 
water,  or  a  solution  of  a  very  dilute  nature,  water  passed  from  the  out- 
side into  the  interior,  and  the  protoplasmic  layer  was  distended  so  as 
to  fill  the  corners  completely.  The  distension  of  the  cells  of  droop- 
ing flowers,  when  their  stems  are  placed  in  water,  and  the  consequent 
revival,  is  a  familiar  illustration  of  the  same  sort  of  thing.  All  solu- 
tions which  produced  neither  the  one  effect  nor  the  other  on  a  given 
set  of  plant  cells,  were  named  is-osmotic.  The  osmotic  pressures  of 
their  contents  were  the  same  as  the  pressure  of  the  cell  fluid. 


FIG.  99. 


DISSOCIATION  IN  SOLUTION  329 

Since  the  entrance  of  the  solvent  is  due  to  the  dissolved  substance, 
and  the  solvent  is  really  drawn  forcefully  into  the  solution,  it  might 
be  more  appropriate  to  call  the  force  osmotic  suction.  Whatever 
it  is  named,  however,  it  is  real,  and  its  value  can  be  measured. 

Professor  Crum  Brown  has  devised  an  arrangement  which  exhibits  the  action 
of  a  perfectly  semi-permeable  membrane  very  strikingly  [Lect.  exp.].  A  concen- 
trated solution  of  calcium  nitrate  is  shaken  with  a  small  amount  of  phenol  (carbolic 
acid),  so  as  to  become  saturated  with  the  latter,  and  the  mixture  is  then  poured 
into  a  tall,  narrow  cylinder.  The  phenol  rises  and  floats  upon  the  surface  of  the 
calcium  nitrate.  The  amount  of  phenol  should  not  be  more  than  sufficient  to 
saturate  the  liquid  and  give  a  layer  a  few  millimeters  in  thickness.  Distilled 
water,  also  saturated  with  phenol,  is  cautiously  introduced  above  alt.  The  water 
on  both  sides  of  the  layer  of  phenol  is  soluble  in  phenol,  and  consequently,  by 
dissolving  in  this  and  passing  out  on  the  other  side,  can  traverse  the  partition. 
The  calcium  nitrate,  however,  which  is  here  the  dissolved  substance,  cannot 
traverse  the  phenol  in  which  it  is  not  soluble.  The  phenol  therefore  constitutes  a 
perfect  semi-permeable  membrane.  If  the  level  of  the  lower  side  of  the  phenol  is 
marked  on  the  outside  of  the  cylinder  by  means  of  a  strip  of  paper,  it  will  be 
found,  as  the  arrangement  is  watched  from  day  to  day,  that  the  water  passes 
through  the  phenol  into  the  solution,  and  the  phenol  rises  higher  and  higher,  until 
finally  it  surmounts  all  the  rest  of  the  liquid. 

The  Phenomena  Logical  Consequences  of  Semi-Permea- 
bility. —  The  passage  of  the  water  into  the  solution  in  which  the 
greater  osmotic  pressure  exists  seems  at  first  paradoxical.  We  must 
remember,  however,  that  the  system,  consisting  of  the  liquids  on  each 
side  of  the  membrane,  can  be  in  equilibrium  only  when  the  activity 
of  the  solvent  on  the  two  sides  is  identical.  But  the  equalization 
of  the  activities  cannot  take  place  by  the  passage  of  part  of  the 
solute  from  one  side  to  the  other,  ^The  membrane  has  been  taken, 
purposely,  of  such  a  nature  that  the  dissolved  substance  is  unable  to 
traverse  it.  The  equalization  must  occur,  therefore,  in  the  only  other 
possible  manner,  namely,  by  the  passage  of  the  solvent  in  the  other 
direction. 

An  imitation  of  this  beftavior  may  easily  be  exhibited  by  the  use  of  gases  [Lect. 
exp.].  A  piece  of  peritoneal  membrane  is  stretched  across  the  mouth  of  a  thistle- 
tube  and  moistened  with  water.  The  tube,  which  has  been  bent  in  U-fofm  to 
serve  as  a  manometer,  contains  a  small  amount  of  some  colored  liquid,  whose 
motions  will  exhibit  any  change  in  pressure  in  the  interior.  When  an  inverted 
cylinder  of  ammonia  gas  is  placed  round  the  head  of  the  thistle-tube,  the  ammonia 
gas  dissolves  in  the  water  on  the  membrane  until  this  water  is  saturated,  that  is, 
until  the  ammonia  molecules  leaving  the  water  are  as  numerous  as  those  entering 


330  INORGANIC  CHEMISTRY 

it.  It  will  be  seen,  however,  that  the  ammonia  solution  really  has  two  surfaces, 
one  of  them  towards  the  interior,  and  the  ammonia  molecules  must  eventually  leave 
both  surfaces  at  the  same  rate  at  which  they  are  landing  upon  one  of  them. 
The  ammonia  gas  being  at  the  pressure  of  the  atmosphere,  the  molecules  of  am- 
monia leaving  the  film  will  produce  a  tension  of  one  atmosphere  of  ammonia  over 
each  surface.  Thus  ammonia  gas  will  be  transferred  from  the  cylinder  to  the 
interior  of  the  thistle-tube  until  its  partial  pressure  in  the  latter  is  equal  to  that  in 
the  former.  The  membrane  is  semi-permeable,  since,  of  the  air  and  ammonia  con- 
tained in  the  thistle-tube,  only  the  ammonia  can  traverse  the  film.  The  con- 
tents of  the  thistle-tube  therefore  correspond  to  the  solution,  air  being  the  solute 
and  ammonia  the  solvent.  The  original  air  in  the  apparatus  was  at  a  pressure  of 
one  atmosphere,  but  the  ammonia,  although  under  no  greater  pressure,  enters 
nevertheless.  Indeed,  it  would  continue  to  do  so  until  the  pressure  inside  became 
equal  to  that  of  the  ammonia  outside  plus  the  original  pressure  of  the  air,  a  total  of 
two  atmospheres.  The  case  corresponds  to  that  of  water  entering  a  solution  whose 
osmotic  pressure  is  one  atmosphere.  It  enters  until  the  contents  of  the  apparatus 
are  under  a  pressure  one  atmosphere  greater  than  that  existing  outside. 

This  experiment  illustrates  the  passage  of  a  substance  into  a  region  of  higher 
pressure,  but  must  not  be  held  to  afford  an  explanation  of  how  osmotic  pressure 
operates.  Osmotic  pressure  cannot  be  explained  as  due  to  impacts  of  water 
molecules  alone  outside  and  to  impacts  of  water  and  solute  molecules  together 
inside,  the  impacts  of  the  latter  constituting  the  excess  of  pressure  inside  (see 
p.  331). 

Measurement  of  Osmotic  Pressure.  —  It  will  be  seen  that 
the  whole  phenomenon  rests  upon  the  fact  that  the  membrane  used 
is  permeable  by  one  of  the  components  only.  The  preparation  of  a 
vessel  of  sufficient  strength,  and  possessing  walls  with  the  maximum 
permeability  by  water  and  the  minimum  permeability  by  dissolved 
substances,  presents  great  difficulties.  A  device  of  Pfeffer's  is  still 
found  to  be  the  best.  A  cylinder  of  porous  porcelain,  much  like  a 
Pasteur  filter-tube,  is  treated  so  that  its  pores  are  partially  filled  with 
a  gelatinous  precipitate  of  cupric  ferrocyanide  (q.v.). 

The  porous  cylinder,  after  removal  under  the  air-pump  of  the  ah*  which  its 
walls  contain,  is  placed  in  a  solution  of  cupric  sulphate.  Its  interior  is  then  filled 
with  a  solution  of  potassium  ferrocyanide.  When  these  two  liquids  meet  by  diffu- 
sion inside  the  wall,  they  interact,  producing  a  dense  precipitate  of  the  substance 
above  mentioned: 

2CuS04  +  K4Fe(CN)6  -»  Cu2Fe(CN)6  J 


The  best  membranes  are  obtained  by  using  a  current  of  electricity  to  cause  the 
copper  to  move  towards  the  cell  on  one  side  and  the  ferrocyanide  radicals  to  move 
towards  it  from  the  other  side. 


DISSOCIATION  IN  SOLUTION 


331 


If  such  a  prepared  vessel,  after  being  filled  with  a  one  per  cent 
sugar  solution,  could  be  closed  by  a  piston  (e.g.,  Fig.  54)  and  be  placed 
in  pure  water,  it  would  be  found  necessary  to  place  weights  on  the 
piston  to  prevent  an  upward  movement,  due  to  access  of  water  to  the 
interior  through  the  walls.  Finally  a  weight  would  be  found  that 
would  just  balance  the  inward  tendency  of  the  water.  With  more 
weight  than  this,  water  would  be  squeezed  out  through  the  pores; 
with  less,  the  water  would  force  its  way  in  and 
the  piston  would  rise.  When  this  weight  has 
been  placed  in  position,  the  water  inside  and 
outside,  having  reached  a  condition  of  equilib- 
rium, must  be  exerting  equal  pressures  on  each 
side  of  the  wall  of  the  vessel.  Hence,  the  ex- 
cess of  pressure  inside  must  be  due  to  the  os- 
motic pressure  of  the  solution.  The  weight 
balancing  the  osmotic  pressure  at  15°,  in  the 
case  of  a  one  per  cent  sugar  solution,  is  found 
to  be  about  0.76  kg.  for  every  sq.  cm.  of  the 
exposed  surface.  Since  1.03  kg.  per  sq.  cm. 
equals  760  mm.,  this  would  indicate  a  pressure 
of  760  X  0.76  -^  1.03,  or  572  mm.  (0.75  at- 
mospheres) . 

In  practice  a  small  bent  tube  opening 
into  the  cylinder  is  used  as  a  manometer 
(Fig.  100).  The  other  end  of  the  tube  is 
closed,  and  some  nitrogen  is  confined  in  this 
end  by  mercury.  The  diminution  in  the  vol- 
ume of  the  nitrogen  registers  the  pressure. 
The  smaller  tube,  drawn  out  to  a  point,  is 
used  for  filling  the  cell  with  the  solution  and 
is  then  sealed  before  the  blow-pipe.  The 
whole  apparatus  is  immersed  in  a  large  bath 
of  water  whose  temperature  can  be  maintained  constant  during  the 
experiment.  Concordant  readings  are  hard  to  get  in  consequence 
of  difficulties  inherent  in  the  preparation  and  use  of  the  ap- 
paratus. 

The  Exact  Relations  of  Osmotic  Pressure.  —  The  value  of 
the  Observed  osmotic  pressure  increases  with  the  concentration  of 
the  solution.  The  chief  relation  is  that  the  osmotic  pressure  is  pro- 
portional to  the  logarithm  of  the  fraction  of  all  the  molecules  in  the 


Fia.  100. 


332  INORGANIC  CHEMISTRY 

solution  which  are  molecules  of  the  solvent.*  This  holds  strictly, 
however,  only  when  there  is  no  chemical  interaction  between  solvent 
and  solute,  when  there  is  no  change  in  volume  consequent  upon 
solution,  and  when  no  heat  change  occurs  upon  dilution  of  the  solu- 
tion. In  actual  fact  many  substances  form  hydrates  when  dissolved 
in  water.  Thus  sugar  seems  to  form  a  penta-  or  a  hexahydrate,  and 
this  removes  a  part  of  the  solvent.  Again,  changes  in  volume  upon 
solution  are  often  considerable  (p.  201).  Also  there  is  always  some 
heat  of  solution  and  often  (cf.  p.  203)  the  value  is  very  great.  Finally, 
it  is  impossible  in  the  case  of  associated  liquids  like  water  (cf.  p.  202) 
to  tell  how  many  molecules  of  the  solvent,  relatively  to  the  number 
of  molecules  of  the  solute,  are  present,  because  we  do  not  yet  know 
how  many  are  H2O,  and  how  many  (H20)2  and  (H2O)3.  Further- 
more, addition  of  a  solute  displaces  the  equilibrium  and  alters  the 
proportions  of  these  numbers.  It  is  thus  impossible  accurately  to 
predict  the  osmotic  pressure  of  a  given  solution,  or  even  accurately 
to  calculate  the  osmotic  pressure  at  one  concentration  from  an 
observation  made  at  another  concentration.  That  van't  Hoff's  gas 
analogy  applies  to  infinitely  dilute  solutions,  but  is  wholly  inappli- 
cable to  ordinary  solutions,  is  easily  shown.  At  0°,  a  0.53  molar 
solution  of  cane  sugar  gives  an  observed  osmotic  pressure  of  13.95 
atmospheres  while  calculation  from  van't  Hoff's  theory  yields  11.79 
atmos.,  and  a  2.2  molar  solution  gives  by  observation  133.74  atmos. 
and  by  calculation  only  49.15  atmos. 

Modes  of  calculation  which,  so  far  as  possible,  take  into  account 
the  modifying  factors  mentioned  above,  cannot  be  discussed  here. 
It  is  sufficient  to  say  that  dissolving  any  substance  in  water,  or  some 
other  solvent,  reduces  the  physical  activity  of  the  solvent.  Thus 
pure  water  is  more  active  than  water  in  a  solution,  and  forces  its 
way  through  a  suitable  membrane  into  the  solution.  Pure  water 
also  shows  a  higher  vapor  pressure,  and  therefore  a  lower  boiling- 
point  (see  p.  337). 

Approximate  Relations  of  Osmotic  Pressure.  —  As  an  aid 

to  memory,  and  as  a  very  rough  indication  of  the  facts,  the  following 
statements,  which  are  approximately  true  for  very  dilute  solutions, 
may  be  made. 

The  osmotic  pressure  is  proportional  to  the  concentration  (par- 
allel of  Boyle's  law).  Thus  the  values  for  sugar  (15°)  are:  0.1  molar 
2.54  atmos.,  0.2  molar  4.99  atmos.,  0.4  molar  9.95  atmos. 

*  For  small  concentrations  the  osmotic  pressure  may  be  taken  without  seri- 
ous error  as  proportional  to  the  molar  fraction  of  the  soluta 


DISSOCIATION  IN  SOLUTION  333 

The  osmotic  pressure  increases  in  proportion  to  the  absolute  tem- 
perature (parallel  of  Charles'  law).  Thus  a  0.1  molar  solution  of 
sugar  gives  at  5°  2.45  atmospheres  pressure,  and  at  50°  2.64.  A 
gas  which  at  278°  Abs.  gave  2.45  atmos.  pressure,  would  give  2.84 
atmos.  at  323°  Abs.,  so  that  the  rule  is  here  over  7  per  cent  in  error. 

Finally,  the  osmotic  pressure  caused  by  a  substance  in  very  dilute 
solution  is  identical  in  value  with  the  gaseous  pressure  which  it  would 
exhibit  if  the  same  quantity  of  it  were  contained  as  a  gas  in  the  same 
volume  at  the  same  temperature.  For  example,  44  g.  of  carbon 
dioxide  in  the  gaseous  condition  fills  the  G.M.V.  (22.4  1.),  and  at  0° 
exercises  a  pressure  of  one  atmosphere.  When  we  dissolve  the  same 
quantity  of  the  same  substance  in  22.4  1.  of  any  solvent  at  the  same 
temperature,  it  causes  approximately  one  atmosphere  of  osmotic 
pressure. 

These  facts  apply  to  substances  which  are  not  acids,  bases,  nor 
salts.  We  shall  learn  in  the  next  section  that  the  osmotic  pressures 
of  the  members  of  these  three  classes  of  substance  are  frequently 
abnormally  high,  but  that  the  abnormality  is  easily  explained. 

Osmotic  pressure  (or  suction)  is  a  subject  of  great  interest  in 
connection  with  the  physiology  of  plants  and  animals.  The  revival 
of  a  withered  flower  has  been  mentioned  (p.  328).  Similarly,  the 
ascent  of  the  water  from  the  soil  into  the  roots  and  through  the 
stem  of  a  growing  plant  is  explained.  In  the  animal  body  also, 
osmosis  plays  a  large  part. 

Osmotic  Pressure  and  Dissociation  in  Solutions.  —  What 
inference  is  to  be  drawn  in  the  cases  in  which  abnormally  high  osmotic 
pressures  are  observed?  In  view  of  the  fact  that  the  pressure  depends 
on  the  fraction  of  foreign  particles  (molecules)  in  the  given  volume, 
we  must  infer  that  where  the  pressure  is  greater,  more  foreign  par- 
ticles are  present  in  the  given  volume  than  we  had  supposed.  In 
other  words,  dissociation  of  the  original  molecules  must  have 
occurred.  This  phenomenon  is  observed  whenever  acids,  bases,  or 
salts  in  aqueous  solution  are  under  observation.  Thus  a  solution 
of  sugar,  which  does  not  belong  to  these  classes,  containing  342  g. 
in  the  G.M.V. ,  exhibits  approximately  the  normal  osmotic  pressure 
of  one  atmosphere  at  0°.  A  solution  of  one  molecular  weight  <rf 
potassium  chloride  (74.5  g.)  in  the  same  volume  of  water,  however, 
exhibits  an  osmotic  pressure  of  about  1.88  atmospheres  at  0°.  The 
greater  pressure  must  be  due  to  the  fact  that,  although  the  number 
of  molecules  of  potassium  chloride  taken  is  the  same  as  in  the  case 


334 


INORGANIC  CHEMISTRY 


of  sugar,  the  number  of  actual  particles  is  greater,  —  is,  in  fact,  88 
per  cent  greater.  Now  the  multiplication  of  particles  from  potas- 
sium chloride  molecules  can  occur  only  by  their  dissociation  into 
particles  of  K  and  Cl  by  a  chemical  change  represented  by  the 
equation  KC1  <=*  K  -f-  Cl.  In  this  case,  seeing  that  each  original 
molecule  can  give  but  two  particles,  the  excess  of  pressure  indi- 
cates that  0.88  (88  per  cent)  of  the  molecules  of  potassium  chloride 
have  been  broken  up.  Comparison  shows  that  the  degree  of  disso- 
ciation for  equi-molar  solutions  of  different  acids,  bases,  or  salts 
varies  widely.  For  the  same  substance,  it  is  always  relatively 
greater  in  dilute  than  in  concentrated  solutions. 

It  will  be  seen  that  we  have  thus  a  purely  physical  and  perfectly 
independent  confirmation  of  the  indications  already  found  in  the 
chemical  behavior  of  substances  of  this  kind.  In 
practice,  on  account  of  the  experimental  difficulties, 
this  method  is  not  used  for  measuring  the  degree 
of  dissociation. 

DEPRESSION  IN  THE  FREEZING-POINT 
OF  A  SOLVENT 

Measurement    of    Freezing-Points.  —  The 

task  consists  in  measuring  exactly  the  temperature 
at  which  a  previously  weighed  quantity  of  the 
solvent  freezes,  and  then,  after  dissolving  in  it  a 
known  weight  of  some  soluble  substance,  deter- 
mining the  freezing-point  once  more.  The  abso- 
lute values  of  these  two  points  are  not  required, 
it  is  simply  the  difference  between  them  that  has 
to  be  known  with  exactness  (cf.  p.  200).  By  means 
of  a  very  delicate  thermometer  (Fig.  101)  having 
only  six  degrees  on  the  whole  scale,  the  tempera- 
ture of  the  freezing  liquid  may  be  read  to  one  one- 
thousandth  of  a  degree.  A  reservoir  at  the  top 
enables  us  to  add  to,  or  subtract  from,  the  mer- 
cury contained  in  the  bulb  and  column,  and  so 
the  same  instrument  may  be  used  with  solvents 
having  widely  different  freezing-points.  When 
water  is  being  employed  as  the  solvent,  the  outer 
jar  must  be  filled  with  a  freezing  mixture  of  ice  and  water  contain- 
ing salt.  With  solutions  in  benzene,  ice  and  water  are  used  alone. 
To  avoid  super-cooling  the  solvent  or  solution  must  be  vigorously 


FIG.  101. 


DISSOCIATION  IN  SOLUTION  335 

stirred  after  it  has  been  cooled  down  to  a  point  just  below  the 
freezing-point. 

Laws  of  Freezing-Point  Depression.  —  The  depression  is 
directly  proportional  to  the  weight  of  dissolved  substance  in  a  given 
amount  of  the  solvent.  The  depression  is  inversely  proportional 
to  the  amount  of  solvent.  Thus,  if  we  double  the  concentration 
of  the  solution,  the  depression  in  the  freezing-point  is  doubled. 
Further,  equal  numbers  of  molecules  of  different  solutes  in  the  same 
quantity  of  solvent  give  equal  depressions.  Or,  in  other  words,  the 
depression  is  proportional  to  the  concentration  of  the  molecules  of 
the  solute.  Thus,  solutions  containing  342  g.  of  sugar 
or  46  g.  of  alcohol  (C2H6O),  or  74  g.  of  methyl  acetate 
in  1000  g.  of  water,  show  a  depression  below  the  freezing-point  of 
water  of  1.86°  in  each  case,  that  is,  such  solutions  will  freeze  close  to 
—  1.86°.  This  depression  produced  by  a  mole  of  the  solute  in  1  1. 
of  water  is  called  the  molecular  depression  constant  and  has  a  dif- 
ferent value  for  each  solvent.  For  solutions  of  the  same  molecular 
concentration  in  benzene  (f.-p.  5.48°)  the  depression  is  5°,  in 
phenol  (carbolic  acid)  7.3°.  Combining  these  facts  in  one  expression: 

The  observed  depres-  )  m  of  Solute  1000 


sion  in  an  aqueous  \  =1.86  X._       '  X 


solution 


Mol.  Wt.  of  Solute    Wt.  of  Solvent 


For  other  solvents,  the  corresponding  value  of  the  depression  constant 
is  substituted  for  1.86°. 

These  principles  may  be  expressed  mathematically  in  a  form  which  is  con- 
venient for  use.  If  A  represent  the  depression  in  any  actual  experiment,  5  the 
depression  produced  by  one  molecular  weight  in  1000  grams  of  solvent,  W  the 
weight  of  the  substance,  M  its  molecular  weight,  and  g  the  weight  of  the  solvent  in 
grams,  then: 

W  X  1000 


In  the  case  of  water,  as  we  have  seen,  5  is  1.86°.     For  each  solvent  the  value  of  8 
must  be  determined  by  means  of  a  substance  of  known  molecular  weight. 

These  laws  describe  the  facts  most  exactly  when  the  solutions  are 
dilute.  They  hold  only  when  there  is  no  chemical  interaction 
between  solute  and  solvent,  and  when  the  crystals  frozen  out  consist 
of  the  pure  solvent.  If  the  crystals  contain  the  same  proportion  of 
the  solute  as  does  the  solution,  no  depression  is  observed;  if  they 


336  INORGANIC  CHEMISTRY 

contain  more,  an  elevation  in  the  freezing  point  is  noted.  Even  so, 
however,  acids,  bases,  and  salts  dissolved  in  water  present  many 
apparent  exceptions  and  must  be  discussed  separately. 

Determination  of  Molecular  Weights.  —  When  the  depres- 
sion constant  of  a  solvent  has  once  been  ascertained  by  means  of  a 
substance  of  known  molecular  weight,  this  method  may  be  used  for 
determining  the  molecular  weight  of  other  substances  which  are 
soluble  in  the  same  liquid.  All  the  other  factors  can  be  observed 
and  substituted  in  the  formula.  This  method  is  especially  useful 
when  the  substance  cannot  be  converted  into  vapor  without  under- 
going decomposition  (see  Hydrogen  peroxide,  p.  318). 

Abnormal  Freezing -Point  Depression:  Dissociation  in  So- 
lution. —  The  substances  which  present  the  most  conspicuous  ex- 
ceptions to  the  above  rules  are  acids,  bases,  and  salts  in  aqueous 
solution.  With  most  of  these,  the  depression  produced  is  greater 
than  we  should  expect  from  the  concentration  of  the  solution.  Thus, 
in  an  actual  experiment,  two  equi-molar  solutions  were  compared. 
One  contained  one  mole  (74  g.)  of  methyl  acetate,  and  the  other 
one  mole  (58.5  g.)  of  sodium  chloride,  each  dissolved  in  2000  g.  (2 
liters)  of  water.  The  freezing-points  observed,  on  the  arbitrary 
scale  of  the  thermometer,  were: 

Pure  water 3.580°        Pure  water      .  >!t!  :1    .     3.580° 

Solution  of  methyl  acetate  .     2.610°         Solution  of  salt  .    .    .    .     1.902° 


Depression 0.970°        Depression      1.678° 

0.970° 
Excess  depression  by  salt  0 . 708° 

The  solution  of  methyl  acetate,  as  it  contained  only  0.5  moles  of 
the  solute  per  liter  of  water,  showed,  as  it  should  do,  about  half  the 
average  molecular  depression  (1.86°,  p.  335).  This  is  typical  of  the 
class  of  substances  showing  normal  behavior.  Sugar,  alcohol,  and 
hundreds  of  other  substances,  in  solutions  of  the  same  molar  con- 
centration, would  have  given  the  same  value. 

The  freezing-point  of  the  salt  solution,  however,  was  much  lower. 
If  this  solution  had  contained  the  same  concentration  of  dissolved 
particles  as  the  other  solution,  its  depression  would  have  been  0.970° 
likewise.  The  number  of  particles  must  therefore  have  been  greater 
than  we  should  have  expected  from  the  number  of  molecules  taken. 
In  other  words,  a  portion  of  the  molecules  of  the  salt  must  have  been 


DISSOCIATION   IN  SOLUTION  337 

broken  up,  and  the  excess  depression,  0.708°,  must  have  been  due  to 
the  extra  particles  produced  by  dissociation.  Now  sodium  chloride 
molecules  cannot  give  more  than  two  particles  each,  and  the  depres- 
sion is  proportional  to  the  number  of  particles.  It  follows,  therefore, 
that  -J^f ,  or  0.732  (73.2  per  cent)  of  the  molecules  were  dissociated: 

(27%)  NaCl  <=>  (Na)  +  (Cl)  (73%). 

This  result  is  typical  also.  Acids,  bases,  and  salts  of  which  one 
mole  is  dissolved  in  two  liters  of  water,  are  found  to  give  irregular 
values,  all  more  or  less  in  excess  of  0.970°.  Those  which  contain 
but  two  radicals,  like  sodium  chloride  (NaCl).  and  potassium  nitrate 
KN03,  give  values  between  0.970°  and  2  X  0.970°.  Substances  like 
calcium  chloride  CaCl2  and  sodium  sulphate  N^SCX  give  depressions 
approaching  three  times  the  normal  value:  their  molecules  contain 
three  radicals.  The  excess  depression  depends,  therefore,  upon 
the  number  of  particles  which  each  molecule  can  furnish,  and  upon 
the  proportion  of  all  the  molecules  which  is  dissociated  into  these 
fragments. 

In  the  case  of  an  acid,  base,  or  salt,  the  depression  is  not  strictly 
proportional  to  the  concentration.  Thus,  one  mole  of  salt  in  four 
liters  of  water  does  not  give  half  the  depression  of  the  two-liter 
solution  (0.839°)  but  somewhat  more  (about  0.844°).  The  same 
method  of  calculation  indicates,  therefore,  a  greater  degree  of  dissocia- 
tion (about  79  per  cent)  in  the  more  dilute  solution.  This  dissocia- 
tion, is,  therefore,  a  reversible  chemical  change. 

Acids,  bases,  and  salts,  so  far  as  they  are  soluble  in  materials  like 
toluene,  benzene,  chloroform,  and  carbon  bisulphide,  exhibit  simply 
normal  depressions  in  these  solvents.  It  appears,  therefore,  that 
dissociation  does  not  take  place  in  many  solvents.  In  common 
experience  it  is  encountered  only  in  solutions  in  water,  and  in  alcohol. 

Abnormal  Boiling-Point  Elevation.  —  If  space  permitted,  a 
series  of  statements  might  be  made  in  regard  to  the  boiling-points  of 
solutions  (cf.  p.  198)  which  would  be  closely  parallel  to  those  about 
freezing-points.  The  boiling-point,  as  we  have  seen,  is  elevated, 
however,  by  the  introduction  of  a  soluble  body.  Thus,  when  water 
is  the  solvent,  one  mole  of  a  solute  in  1000  g.  of  the  solvent  normally 
raises  the  boiling-point  0.52°  (that  is,  from  100°  to  100.52°).  One 
molecular  weight  of  sodium  chloride  (58.5  g.),  however,  will  elevate 
the  boiling-point  of  the  water  0.87°  instead  of  0.52°.  The  effect  is 
0.35°,  or  67  per  cent  greater,  indicating  dissociation  of  this  pro- 


338  INORGANIC  CHEMISTRY 

portion  of  the  NaCl  molecules.  In  more  dilute  solutions,  the  eleva- 
tion is  relatively  greater.  Salts  containing  more  than  two  radicals, 
like  Ca(Cl)2,  give  elevations  of  more  than  twice  the  normal  value. 
In  solvents  like  benzene  and  carbon  disulphide,  however,  no  abnor- 
mally large  elevation  is  observed  with  any  solute.  The  phenom- 
ena are,  in  fact,  parallel  with  those  connected  with  the  freezing-point. 

Comparison  of  the  Results  of  the  Three  Methods.  —  When 
we  measure  the  osmotic  pressure,  the  freezing-point  depression,  and 
the  elevation  in  the  boiling-point  of  the  same  solution,  and  calculate 
the  degree  of  dissociation  from  the  result  of  each  measurement,  we 
find  that  the  values  obtained  are  usually  identical,  within  the  limits 
of  error  to  which  the  methods  are  liable.  Indeed,  the  theory  of  this 
subject  enables  us  to  connect  the  osmotic  pressure  by  a  mathematical 
relation  with  the  other  two  phenomena,  and  to  calculate  any  one  of 
the  three  from  any  other.  Thus  the  indications  of  dissociation 
found  in  the  chemical  behavior  of  acids,  bases,  and  salts  (p.  326) 
are  fully  confirmed  by  a  study  of  the  physical  properties  of  their 
solutions.* 

The  connection  between  the  three  sets  of  phenomena  cannot  be  explained 
here.  It  is  treated  in  all  works  on  Physical  Chemistry.  It  may  be  pointed  out, 
however,  that,  in  one  essential  respect,  experiments  in  osmotic  pressure,  and  in 
the  freezing  and  boiling  of  solutions,  are  all  alike.  The  perception  of  osmotic 
pressure  involves  a  partition  which  the  solvent  alone  can  pass,  and  the  osmotic 
pressure  for  a  given  solution  is  the  one  required  to  force  the  solvent  out.  In 
freezing  a  solution,  pure  ice  is  separated,  and  so  a  similar  extrusion  of  a  part  of  the 
pure  solvent  is  effected.  In  a  boiling  solution,  for  which  the  above  rules  hold,  the 
vapor  is  composed  of  the  pure  solvent,  and  the  solute  remains  behind.  The  rela- 
tion between  the  three  operations  lies  in  the  fact  that  in  each  case  the  same  thing, 
namely,  the  separation  of  a  part  of  the  solvent,  is  done.  Each  method  effects  this 
in  a  different  way.  But  the  expressions  representing  the  work  done,  in  terms  of 
the  factors  which  define  the  work  in  each  case,  can  be  equated  in  pairs  and  the 
required  relation  established.  Thus  the  molecular  depression  of  the  freezing- 
point,  or  the  molecular  elevation  in  the  boiling-point,  as  we  have  defined  them,  is 
equal  to  0.002  T2  -T-  q,  where  T  is  the  absolute  temperature  of  the  freezing-  or 
boiling-point,  and  q  is  the  heat  of  fusion  or  vaporization,  as  the  case  may  be. 
Water,  for  example,  freezes  at  273°  abs.,  and  its  heat  of  fusion  is  79  cal.  per  gram, 
from  which  the  calculated  molecular  depression,  0.002  X  2732  -T-  79,  or  1.88°,  is 

*  Recent  observations,  showing  that  in  some  cases  rapid  double  decomposi- 
tions of  the  normal  kind  take  place  in  solutions  which  exhibit  no  physical  evidence 
of  the  existence  of  dissociation,  demonstrate  that  it  would  have  been  unsafe  to 
infer  dissociation  from  chemical  evidence  alone. 


DISSOCIATION  IN  SOLUTION  339 

obtained.  Similarly,  using  the  boiling-point,  373°  abs.,  and  the  heat  of  vaporiza- 
tion of  water,  540  cal.  per  gram,  we  calculate  the  molecular  elevation  of  the 
boiling-point  to  be  0.518°. 

It  ought  to  be  added  that  abnormally  smatt  osmotic  pressures,  freezing-point  de- 
pressions, and  boiling-point  elevations,  are  also  frequently  observed.  This  occurs, 
however,  almost  wholly  in  non-aqueous  solvents,  such  as  benzene.  It  is  shown 
particularly  by  substances  containing  oxygen,  and  is  even  noticed  in  the  case  of 
acids,  bases,  and  salts.  By  parity  of  reasoning  we  infer  that  in  these  cases  associa- 
tion (cf.  p.  282)  of  the  molecules  has  occurred,  and  that  the  physical  unit  of  the 
solute  in  these  solvents  is  larger  than  the  ordinary  molecule. 

THE  APPLICATION  OF  THESE  CONCLUSIONS  IN  CHEMISTRY 

The  Constitution  of  Solutions  of  Acids,  Bases,  and  Salts. 

—  The  composition  of  solutions  which  are  normal  or  abnormal,  in 
respect  to  osmotic  pressure,  freezing-point,  and  boiling-point,  may  be 
shown  thus: 


SOLUTES. 

DISSOLVED    IN    WA- 
TER, .ALCOHOL,    ETC. 

DISSOLVED  IN 
TOLUENE,    CHLORO- 
FORM, ETC. 

Acids,  bases,  salts  

Abnormal 

Normal 

Other  substances    

Normal 

Normal 

It  appears  that  water  and  some  other  solvents  have  the  power  of 
decomposing  acids,  bases,  and  salts.  Such  solvents  have,  in  fact, 
an  effect  on  these  materials  that  resembles,  outwardly  at  least,  the 
effect  which  heat  has  on  many  substances  (cf.  p.  260),  they  cause 
dissociation: 

CaCl2<=±(Ca) 


In  consequence  of  this,  our  view  of  the  nature  of  an  aqueous 
solution  of  hydrogen  chloride  HC1,  or  common  salt  NaCl,  or  sodium 
hydroxide  NaOH,  or  any  of  the  substances  of  the  classes  which  these 
represent,  may  now  be  stated  in  definite  terms.  Such  a  solution 
contains,  besides  undivided  molecules  of  the  solute,  at  least  two  other 
kinds  of  material,  H,  Na,*  Cl,  OH,  etc.,  which  result  from  the  break- 
ing up  of  the  molecules.  We  shall  see  that  these  subdivisions  of  the 
original  molecules  have  distinct  physical  and  chemical  properties  of 
their  own.  The  descriptions  of  the  "properties"  of  the  solutions,  as 

*  The  objection  that  separate  atoms  of  sodium  could  not  remain  free  in  water, 
will  be  disposed  of  later, 


340  INORGANIC  CHEMISTRY 

they  used  to  be  given  in  chemistry,  were  really  a  confused  statement 
of  the  properties  of  the  different  components  of  a  mixture. 

The  suggestion  that  the  multiplication  of  particles  takes  place  by  interaction 
of  the  salt  with  part  of  the  water,  NaCl  +  H2O  <=±  NaOH  -f  HC1,  resulting  in  the 
production  of  two  molecules  of  dissolved  matter  from  one,  is  open  to  several  fatal 
objections.  In  the  case  of  a  highly  dissociated  salt,  according  to  this  explanation, 
the  mixing  of  the  acid  and  base  in  dilute  solution  should  result  in  no  particular 
change  and  give  rise,  therefore,  to  no  development  of  heat.  But  the  heat  of 
neutralization  is  very  great  in  such  cases.  This  is  an  example  of  a  stochastic 
hypothesis  (p.  176),  be  it  noted,  and  its  verity  or  falsity  can  be  put  to  the  test  at 
once.  Its  inapplicability  is  further  seen  in  the  fact  that  it  cannot  explain  the  dis- 
sociation of  acids  and  bases  themselves. 

The  free  radicals,  of  whose  existence  we  have  thus  become  con- 
vinced, constitute  a  new  set  of  materials  (with  appropriate  names. 
See  p.  356).  Thus  the  hydrogen  radical  of  acids,  although  a  form  of 
uncombined  hydrogen,  differs  totally  from  the  gas  which  is  com- 
posed of  the  same  material.  The  gas  has  no  sour  taste  or  effect  upon 
litmus;  these  are  properties  of  the  free  radical.  The  gas  is  very 
slightly  soluble  in  water,  while  the  hydrogen  radical  exists  as  a 
separate  substance  only  in  solution.  Again,  substances  with  the 
composition  of  the  radicals  N03  and  SO4  are  not  known  at  all  except 
in  solutions.  The  chief  peculiarity  of  these  substances  is  that  a 
solution  cannot  be  made  which  contains  less  than  two  kinds  of  them 
side  by  side. 

Exercises.  —  1.  What  gaseous  pressure  would  be  exerted  by  a 
gas  of  the  same  molecular  concentration  as  a  one  per  cent  solution  of 
sugar  at  15°  (p.  331)?  Compare  the  answer  with  the  osmotic  pressure 
of  the  solution. 

2.  What  depression  in  the  f.-p.  of  water  will  be  produced  by 
dissolving  10  g.  of  bromine  in  1  kg.  of  this  solvent? 

3.  What  depressions  in  the  f.-p.  of  benzene  and  of  phenol  would 
be  produced  by  10  g.  of  bromine  to  1  kg.  of  the  solvent,  if  no  chemical 
action  took  place? 

4.  What  is  the  molecular  depression-constant  of  a  solvent  in 
which  5  g.  of  iodine  in  500  g.  of  the  solvent  lowers  the  f.-p.  0.7°? 

5.  What  is  the  degree  of  dissociation  of  zinc  sulphate  if  5  g.  of  it 
dissolved  in  125  g.  of  water  produce  a  lowering  of  0.603°  in  the  f.-p.? 
What  is  the  molecular  concentration  of  each  of  the  three  substances 
present  in  this  solution? 


DISSOCIATION  IN  SOLUTION  341 

6.  What  will  be  the  approximate  b.-p.  of  a  solution  of  common 
salt,  saturated  at  100°  (p.  191)?    Assume  that  the  solute  is  50  per 
cent  dissociated. 

7.  In  a  decinormal  solution,  potassium  chloride  is  86  per  cent 
dissociated.     What  is  the  freezing-point  of  this  solution? 

8.  If  5  g.  of  a  substance,  dissolved  in  1000  c.c.  of  water,  give 
a  solution  freezing  at  —0.2°,  what  is  the  molecular  weight  of  the 
substance? 

9.  6  g.  of  a  substance  when  'dissolved  in  200  c.c.  of  water  give 
a  boiling-point  of  102.6°.     What  is  the  molecular  weight  of  the 
substance? 

10.  1.6  g.  of  naphthalene  CioHg  when  dissolved  in  25  g.  of  benzene 
(freezing-point  5.48°)  gives  a  solution  which  freezes  at  3.03°.     When 
2.44  g.  of  another  substance  are  dissolved  in  the  same  amount  of 
benzene,  the  solution  freezes  at  3.52°.     What  is  the  molecular  weight 
of  the  latter  substance? 

11.  The  elevation  of  the  boiling-point  in  the  above  solution  of 
naphthalene  is  1.285°.     What  elevation  of  the  boiling-point  is  pro- 
duced in  the  second  solution? 


CHAPTER  XVIII 
IONIZATION 

Introductory.  —  As  we  have  seen,  acids,  bases,  and  salts,  when 
dissolved  in  water,  interact  with  one  another  by  interchanging  radicals 
(p.  324).  We  have  also  learned  that  the  same  solutions  have  ab- 
normal values  for  their  freezing-points,  boiling-points,  and  osmotic 
pressures.  These  facts  indicate  dissociation  into  the  radicals  (p. 
339).  Now  precisely  these  solutions  have  a  property  which  is  not 
shared  by  any  other  solutions,  namely,  that  of  being  conductors  of 
electricity  and  suffering  chemical  decomposition  by  the  passage  of  the 
current.  Such  solutions  are  called,  in  consequence,  electrolytes, 
and  the  process  is  named  electrolysis.  Now  the  natural  inference 
from  the  foregoing  facts  is  that  the  electricity  is  carried  by  the 
liberated  radicals.  Our  first  aim  in  the  present  chapter  is  to  show, 
by  a  study  of  the  chemical  changes  taking  place  in  electrolysis,  that 
this  inference  is  correct.  We  then  proceed  to  discuss  the  nature 
of  ions  as  a  land  of  molecules.  Next,  we  devote  ourselves  to  the 
explanation  of  electrolysis,  to  the  equilibrium  between  the  ions  and 
the  remaining,  undissociated  molecules,  and  to  conductivity  phe- 
nomena as  a  means  of  measuring  the  fraction  ionized.  Finally,  we 
deduce  the  ^relation  between  extent  of  ionization  and  chemical 
activity. 

Incidentally,  the  facts  to  be  given  provide  the  means  of  under- 
standing the  electrolytic  processes,  many  of  them  of  great  impor- 
tance in  chemical  industries,  to  which  frequent  reference  is  made  in 
later  chapters. 

Non- Electrolytes.  —  To  clear  the  ground,  we  should  first  note 
the  fact  that  only  solutions  (as  a  rule)  possess  both  of  the  properties 
in  question,  namely  that  of  conducting  and  that  of  being  decom- 
posed by  the  current.  Some  substances,  notably  the  metals  and 
materials  like  carbon,  are  conductors.  But  they  are  not  changed 
chemically  by  the  current.  Again,  single  substances,  even  when 
they  are  such  as,  if  mixed,  yield  electrolytes,  are  not  conductors  at 
ordinary  temperatures.  Thus  hydrogen  chloride,  whether  gaseous 

342 


IONIZATION  343 

or  liquefied,  is  a  nonconductor,  and  water  is  a  very  feeble  conduc- 
tor, although  the  solution  of  the  two  conducts  exceedingly  well. 
Dry  acids,  bases,  and  salts,  except  when  at  a  high  temperature  and 
fused,  are  likewise  nonconductors.  Furthermore,  even  amongst 
solutions,  not  all  are  conductors.  Solutions  of  sugar  and  other 
substances  of  the  same  class  (p.  335),  which  have  normal  freezing- 
points,  are  nonconductors.  Only  solutions  of  acids,  bases,  and 
salts  in  certain  specified  solvents,  of  which  the  commonest  is  water, 
are  electrolytes  at  ordinary  temperatures. 

Chemical  Changes  Taking  Place  at  the  Electrodes  During 
Electrolysis.  —  When  the  wires  from  a  battery  are  attached  to 
platinum  plates  immersed  in  any  electrolyte  (e.g.,  Fig.  87,  p.  216,  or 
Fig.  16,  p.  29),  we  observe  that  the  products  appearing  at  the  two 
electrodes  are  always  different.  They  may  be  of  several  kinds 
physically,  and  will  be  secured  for  examination  variously  according 
to  their  nature.  When  they  are  gases,  which  are  not  too  soluble, 
they  may  be  collected  in  inverted  tubes  filled  with  the  solution. 
Solids,  if  insoluble  in  the  liquid,  will  either  remain  attached  to  the 
electrode  or  fall  to  the  bottom  of  the  vessel  as  precipitates.  .  Soluble 
substances,  on  the  other  hand,  will  usually  not  be  visible.  They 
may  be  handled  by  interposing  a  porous  partition  of  some  description 
which  will  restrain  the  diffusion  of  the  dissolved  body  away  from  the 
neighborhood  of  the  electrode,  while  not  interfering  appreciably  with 
the  passage  of  the  current.  Surrounding  one  electrode  with  a  porous 
battery  jar  is  a  convenient  method  for  effecting  this. 

Of  the  various  illustrations  which  we  have  encountered,  the  elec- 
trolysis of  hydrochloric  acid  (p.  216)  happens  to  have  been  the  only 
one  which  delivered  both  components  of  the  solute  with  a  minimum 
of  modification  at  the  electrodes: 

Neg.  Wire,  H2< H.C1 >C12,  Pos.  Wire. 

Hydrogen  does  not  interact  with  water,  and  chlorine  interacts  very 
incompletely,  so  that  the  molecular  substances  H2  and  Cl2  are 
promptly  formed  from  the  elements  H  and  Cl  which  are  liberated. 
The  chlorides,  bromides,  and  iodides  of  those  metals  which  do  not 
interact  with  water  (p.  129)  give  equally  simple  results: 

Neg.  Wire,  Cu< Cu.Br2 »Br2,  Pos.  Wire. 

Thus  the  solute  seems  to  be  split  into  its  radicals,  and  in  electrolysis, 
the  radicals,  if  they  do  not  interact  with  water,  are  set  free.  A 


344  INORGANIC  CHEMISTRY 

substance  thus  set  free  is  called  a  primary  product  of  the  electrolysis. 
In  the  foregoing  instances  both  products  are  primary. 

Usually  the  chemical  change  is  more  complex.  Thus,  when 
dilute  sulphuric  acid  is  electrolyzed,  hydrogen  and  oxygen  are  lib- 
erated at  the  negative  and  positive  electrodes,  respectively.  But 
these  products  do  not  account  for  the  whole  of  the  constituents 
(H2SO4).  We  therefore  proceed  to  examine  the  materials  in  solution 
round  the  electrodes.  It  is  found  that,  as  the  action  progresses, 
sulphuric  acid  accumulates  round  the  positive  wire,  while  the  liquid 
in  the  neighborhood  of  the  other  pole  is  gradually  depleted  of  this 
substance.  In  view  of  this  fact  we  easily  explain  the  phenomenon. 
Evidently  the  substance  divides  into  its  radicals,  H  and  SO4,  but  the 
SO4  must  interact  with  the  water  to  produce  sulphuric  acid  and 
oxygen:  2S04  +  2H20  ->  2H2S04  +  O2.  The  whole  change  may 
therefore  be  tabulated  as  follows: 

Neg.  Wire,  H2< H2.S04 >02  and  H2SO4,  Pos.  Wire. 

Hence  the  hydrogen  is  a  primary  product,  but  the  oxygen  and  sul- 
phuric acid  are  secondary  products.  All  acids  give  hydrogen  alone 
at  the  negative  electrode,  whatever  may  be  the  product  at  the 
positive. 

If  we  electrolyse  cupric  nitrate  solution,  we  obtain  a  red  deposit 
of  metallic  copper  on  the  negative  plate  and  at  the  positive  end 
oxygen  and  nitric  acid  are  formed.  We  infer,  therefore,  that  the 
division  of  the  original  molecule  was  into  Cu  and  N03,  but  that  the 
latter  interacted  with  the  water:  4NO3  +  2H20 >  4HNO3  +  O2: 

Neg.  Wire,  Cu< Cu.(NO3)2 >O2  and  HN03,  Pos.  Wire. 

With  a  solution  of  potassium  chloride  we  find  hydrogen  and 
chlorine  appearing  at  the  negative  and  positive  electrodes,  respec- 
tively. Litmus  paper,  however,  shows  the  presence  in  the  solution 
of  a  base  (potassium  hydroxide,  KOH)  at  the  negative  end.  We 
infer  that  the  parts  of  the  parent  molecules  are  K  and  Cl.  The 
former,  since  it  resembles  sodium  in  being  much  more  active  than 
hydrogen  (p.  216),  is  more  difficult  to  liberate.  Hence  hydrogen  is 
liberated  instead,  and  potassium  hydroxide  remains  in  the  liquid: 
2K  +  2HOH  ->  2KOH  +  H2: 

Neg.  Wire,  H2  and  KOH<—  -K.C1 >C12,  Pos.  Wire. 

We  are  confirmed  in  this  explanation  when  we  employ  a  solution 


IONIZATION  345 

containing  a  mixture  of  salts  of  copper  and  silver.  The  latter, 
being  the  less  active  metal,  is  first  deposited,  alone.  The  copper 
is  liberated  only  after  all  the  silver  has  been  set  free. 

Having  now  before  us  the  results  of  electrolyzing  some  typical 
substances,  we  bring  these  results  into  relation  with  the  facts  de- 
scribed in  Chapter  XVII.  Acids  contain  hydrogen  which  pos- 
sesses certain  specific  properties  (p.  324),  and  in  electrolysis  all 
acids  divide  so  as  to  give  up  this  constituent  alone  at  one  electrode. 
The  evidence  that  the  other  radical  has  different  electrical  proper- 
ties which  carry  it  to  the  opposite  plate  is  conclusive.  Again,  salts 
undergo  double  decomposition  in  which  they  exchange  radicals 
with  acids,  bases,  and  other  salts  (p.  324),  and  we  find  that  it  is 
these  very  radicals  which  are  withdrawn  from  the  solution  by  the  in- 
fluence of  the  electricity.  Furthermore,  the  radicals  exist  free  in 
the  solution,  being  formed  by  dissociation  of  the  molecules  (p.  339). 
Hence  the  function  of  the  electricity  seems  simply  to  consist  in  sifting 
apart  the  two  kinds  of  free  radicals  which  each  solution  contains. 
It  only  remains  for  us  to  explain  in  detail  the  sifting  action  of  the 
current.  Before  turning  to  this  explanation  of  the  phenomena, 
however,  there  is  one  question  which  may  be  answered  in  passing. 
Since  a  solution  may  eventually  be  cleared  of  all  the  hydrochloric 
acid,  for  example,  which  it  contains,  we  should  like  to  know  how 
the  free  radicals  in  the  center  of  the  cell  reach  the  electrodes. 

Ionic  Migration.  —  To  know  how  the  free  radicals  reach  the 
electrodes,  all  that  is  necessary  is  to  take  a  material,  one  (or  both) 
of  whose  radicals  is  a  colored  substance,  and  watch  the  movement  of 
the  colored  material  as  it  drifts  towards  the  electrode.  Most  salts 
that  give  colored  solutions  are  suitable.  In  dilute  cupric  sulphate 
solution,  for  example,  a  freezing-point  determination  shows  that  the 
depression  has  practically  double  the  normal  value.  In  other  words, 
the  dissociation  into  the  radicals,  CuSO4<=^  (Cu)  +  (SO4),  is  almost 
complete.  Now,  the  blue  color  of  this  solution  cannot  be  due  to  the 
remaining  molecules  of  CuSO/i,  for  anhydrous  cupric  sulphate  is 
colorless.  Nor  is  it  due  to  the  color  of  the  (804)  radicals,  for  dilute 
potassium  sulphate  and  dilute  sulphuric  acid  are  both  colorless.  On 
the  other  hand,  all  cupric  salts,  in  dilute  solution,  have  the  same  tint. 
The  color  is  therefore  that  of  the  free  cupric  radical  (Cu).  In  order 
most  clearly  to  see  the  motion  of  the  cupric  radical,  we  place  the 
cupric  sulphate  solution  in  the  middle  of  the  space  between  the 
electrodes,  and  place  between  it  and  the  latter  a  colorless  conducting 


346 


INORGANIC  CHEMISTRY 


FIG.  102. 


solution.     The  motion  of  the  blue  material  across  the  boundary  may 
then  be  easily  observed. 

The  most  convenient  arrangement  is  to  dissolve  the  cupric  sul- 
phate in  warm  water  containing  about  5  per  cent  of  agar-agar  (a 

gelatine  obtained  in  China  from  cer- 
tain sea-weeds),  and  to  fill  with  this 
mixture  the  lower  part  of  a  U-tube 
(Fig.  102).  The  setting  of  the  jelly 
prevents  subsequent  mixing  of  the 
cupric  sulphate  system  of  materials 
with  the  rest  of  the  filling  of  the  tube, 
and  the  consequent  disappearance  of 
the  boundary.  A  few  grains  of  char- 
coal are  scattered  on  the  surface  of 
the  jelly  to  mark  the  present  limits 
of  the  colored  substance,  and  a  solu- 
tion of  some  colorless  electrolyte,  such 
as  potassium  nitrate,  is  added  on 
each  side.  To  prevent  agitation  of 
the  liquid  by  the  effervescence  at  the 
electrodes,  it  is  well  to  use  agar-agar 
with  the  lower  part  of  the  colorl  jss  liquid  also.  The  whole  is  finally 
placed  in  ice  and  water,  to  prevent  melting  of  the  jelly  by  the  heat 
caused  by  resistance,  and  the  current  is  then  turned  on.  The  agar- 
agar  does  not  offer  any  appreciable  resistance  to  the  motion  of 
the  ions,  and  is  presumed  to  form  a  sort  of  open  network  in  the 
solution. 

After  a  time,  we  observe  that  the  blue  cupric  ions  ascend  above  the 
mark  on  the  negative  and  descend  away  from  it  on  the  positive  side. 
In  each  case  there  is  no  shading  off  in  the  tint.  The  motion  of  the 
whole  aggregate  of  colored  radicals  occurs  in  such  a  way  that,  if  the 
contents  of  the  tube  were  not  held  in  place  by  the  jelly,  we  should 
believe  that  a  gradual  motion  of  the  entire  blue  solution  was  being 
observed.  With  a  current  of  110  volts,  and  a  50-watt  lamp  (one- 
half  ampere)  in  series  with  the  cell,  the  effect  becomes  apparent  in  a 
few  minutes. 

Although  the  (804)  radicals  are  invisible,  we  may  safely  infer  that 
they  are  drifting  towards  the  positive  electrode.  Indeed,  this  can  be 
demonstrated  by  interposing  a  shallow  layer  of  jelly  containing  some 
barium  salt  a  little  distance  above  the  charcoal  layer  on  the  positive 
side.  When  the  (804)  radicals  reach  this,  barium  sulphate  begins  to 


IONIZATION  347 

be  precipitated  and  the  layer  becomes  cloudy.     In  similar  ways  the 
progress  of  other  colorless  ions  may  be  rendered  visible. 

It  appears  therefore  that  electrolysis  is  not  a  local  phenomenon, 
going  on  round  the  electrodes  only,  but  that  the  whole  of  the  products 
of  the  dissociation  of  the  solute  are  set  in  motion.  It  is  on  account  of 
this  remarkable  property  of  traveling  or  migrating  towards  one  or 
other  of  the  electrodes  connected  with  a  battery  that  the  individual 
atoms  (like  Cu),  or  groups  of  atoms  (like  SO4)  have  been  named  ions 
(Gk.  iW,  going).  The  term  was  first  applied  by  Faraday  to  the 
materials  liberated  round  the  electrodes. 

Some  writers  use  the  word  "wandering'  for  migration.  But  wandering 
means  rambling  without  any  certain  course  or  object  in  view,  like  diffusion  of  a 
gas.  The  ions  move  like  a  disciplined  regiment,  with  a  straight  front,  and  exactly 
equal  speed  along  that  front.  Migration  is  the  movement  of  a  population  in  a 
definite  direction,  and  is,  therefore,  the  appropriate  term.  The  error  is  due  to 
ignorance  of  German,  in  which  language  Wanderung  means  migration  and  not 
wandering. 

Relative  Speeds  of  Migration  of  Different  Ions.  —  The  speeds 
of  different  ions  may  readily  be  compared.  The  cupric  ion  moves  at 
the  same  speed  whatever  salt  of  copper  we  employ.  In  fact,  the 
speeds  of  all  ions  are  individual  properties  and  are  independent  of  the 
nature  of  other  ions  that  may  be  present.  The  speeds  of  all  are 
increased  by  using  a  current  of  greater  electromotive  force.  Under 
similar  conditions,  the  relative  speeds  of  most  ions  are  in  the  neighbor- 
hood of  50  or  60,  on  the  scale  commonly  used  in  expressing  ionic 
velocities.  Thus,  we  have  (K)  65.3,  (Cl)  65.9,  (Cu)  49.  The  speed 
of  the  hydrogen  ion  of  acids  is  the  greatest  of  all,  318,  while  that  of 
hydroxide  ion  of  bases  (OH)  comes  next,  being  174.  These  are, 
respectively,  about  five  and  two  and  one-half  times  as  fast  as  any 
other  ions. 

The  actual  speeds  of  these  ions  in  dilute  solutions  at  18°,  when 
driven  by  a  potential  difference  of  1  volt  between  plates  1  cm.  apart, 
expressed  in  cm.  per  hour  is:  (K)  2.05,  (Cl)  2.12,  (Cu)  1.6,  (H)  10.8, 
(OH)  5.6,  (S04)  1.6. 

By  an  experiment  similar  to  the  last,  and  devised  by  A.  A.  Noyes,  the  relative 
speeds  of  different  ions  may  be  demonstrated.  The  U-tube  (Fig.  103,  showing 
the  same  tube  A  before  the  current  starts,  and  B  after  it  has  been  passing  for  some 
time)  is  partly  filled  with  agar-agar' emulsion  containing  potassium  chloride  and 
phenolphthalem  (see  Indicators).  On  the  right  side,  a  drop  of  potassium  hy- 
droxide has  been  added  to  render  the  mixture  pink.  On  the  left,  a  drop  of  hydro- 


348 


INORGANIC  CHEMISTRY 


chloric  acid  is  present,  and  the  mixture  is  colorless.  Above  the  charcoal  layer,  in 
che  right  limb,  a  mixture  of  hydrochloric  acid  and  cupric  chloride,  i.e.,  H+  and 
GU++,  and  in  the  left  limb  potassium  hydroxide  solution,  i.e.,  OH~~,  are  placed. 
The  positive  electrode  is  introduced  on  the  right  and  the  negative  on  the  left. 
The  H  and  Cu  ions  drift  away  from  the  former  down  the  tube  towards  the  latter, 
the  OH  ions  away  from  the  latter  down  the  tube  towards  the  former.  The  motion 


H+  &  CU-H-         OH- 
I  I 

Pink  =  OH~ 


FIG.  103. 

of  the  H  is  marked  by  the  disappearance  of  the  pink  color,  that  of  the  Cu  by  the 
advance  of  a  blue  layer,  that  of  the  OH  by  the  progress  of  a  pink  coloration.*  By 
the  time  the  H  ions  have  been  displaced  5?  cm.,  the  Cu  ions  have  moved  1  cm. 
and  the  OH  about  2|  cm.  These  distances  indicate  their  relative  speeds  of  mi- 
gration. 

The  Nature  of  Ions.  —  That  the  molecdles  of  certain  classes  of 
substances,  although  seemingly  without  chemical  interaction  with 
the  water  in  which  they  are  dissolved,  should  nevertheless  be  de- 
composed by  the  influence  of  the  water,  is  strange,  but  not  inconceiv- 
able. Heating  produces  a  somewhat  similar  effect  on  many  sub- 
stances. There  are  two  peculiarities  to  be  accounted  for: 

How  can  the  production  of  a  conducting  medium  by  mixing  two 
nonconductors  be  imagined  to  take  place?  The  solvent  is  a  non- 
conductor, and  the  ions,  even  if  they  are  composed  of  conducting 

*  Bases,  on  account  of  the  (OH)  they  give,  turn  phenolphthalei'n  solutions 
from  colorless  to  pink;  acids,  on  account  of  the  (H)  they  furnish,  turn  it  from  pink 
tc  colorless  (see  Indicators). 


IONIZATION  349 

material,  are  distributed  through  the  liquid  as  independent  particles 
and  cannot  furnish  a  continuous  medium  for  the  stream  of  electricity. 
This  will  be  clear  when  we  remember  that  although  liquid  mercury  is 
an  excellent  conductor,  mercury  vapor,  composed  as  it  is  of  conduct- 
ing particles,  is  not  a  conductor. 

Again,  the  conducting  power  of  the  solution  is  indissolubly  con- 
nected with  the  fact  that  the  original  molecules  of  the  solute  have  been 
broken  up  by  the  solvent  into  smaller  molecules  containing  one  or 
more  atoms.  Why  should  this  particular  kind  of  sub-molecules  be 
attracted  by  electrically  charged  plates,  which  have  been  lowered  into 
the  solution,  when  molecules  of  dissolved  sugar,  for  example,  are  not 
so  attracted? 

An  answer  to  the  second  question  readily  suggests  itself.  The 
only  bodies  which  we  find  to  be  conspicuously  attracted  by  electrically 
charged  objects  are  bodies  which  are  already  provided  with  electric 
charges  of  their  own.  Thus  we  are  led  to  add  the  idea  that  substances 
which  undergo  dissociation  in  solution  divide  themselves  into"  a  special 
kind  of  electrically  charged  molecules. 

Since  the  solution,  as  a  whole,  has  itself  no  charge,  equal  quantities 
of  positive  and  negative  electricity  must  be  produced: 

+  Cr    NaOH  ^  Na+  +  OH~. 


Wild  as  this  supposition  seems  at  first  sight  to  be,  it  turns  out  that 
no  valid  objection  to  it  can  be  raised.  This  means  that  bivalent 
radicals,  on  dissociation,  will  become  ions  carrying  a  double  charge 
and  trivalent  ions  must  carry  a  triple  charge: 

Cuci2  <=>  GU++  +  2cr,      Cuso4  «±  GU++  +  sor, 

K2SO4  <=*  2K+  +  S04=,  FeCl3  <=*  Fe^  +  3C1~. 

In  these  equations,  the  coefficients  multiply  the  charges  as  well  as 
the  radicals  bearing  the  charges,  and  it  will  be  seen  that  the  num- 
bers of  +  and  —  charges  produced  by  each  dissociation  are  equal. 
Hence,  univalent  ions  all  possess  equal  quantities  of  electricity,  and 
other  ions  bear  quantities  greater  than  this  in  proportion  to  their 
valence.  This  is  an  inevitable  inference  from  the  electrical  neu- 
trality of  all  solutions.  An  ion  is  therefore  an  atom  or  group  of 
atoms  bearing  an  electric  charge. 

To  show  that  this  view  of  the  nature  of  the  ions  is  adequate,  after 
a  section  on  Faraday's  law,  we  shall  next  apply  it  to  the  explanation 
of  the  phenomena  of  electrolysis.  After  that  some  seeming  objections 
will  be  discussed. 


350  INORGANIC  CHEMISTRY 

Faraday's  Law.  —  The  above  conclusion  is  confirmed  by  actual 
measurement.  Quantitative  experiments  in  electrolysis  show  the 
most  perfect  adjustment  in  this  respect.  Thus,  in  a  single  cell,  the 
quantities  of  material  liberated  at  the  two  poles  are  invariably 
chemical  equivalents  of  one  another.  With  hydrochloric  acid,  while 
1.008  g.  of  hydrogen  are  being  liberated  at  one  pole,  35.46  g.  of  chlorine 
are  set  free  in  the  same  time  at  the  other.  But  when  cupric  chloride 
CuCl2  is  substituted,  for  every  35.46  g.  (=  Cl)  of  chlorine  set  free, 
only  31.78  g.  (=  J  Cu  =  |  63.57)  of  copper  are  deposited. 

Again,  the  amount  of  any  one  substance  liberated  is  proportional  to 
the  quantity  of  electricity  which  has  traversed  the  cell.  This  is  the 
first  part  of  Faraday's  law. 

Finally,  the  passage  of  equal  quantities  of  electricity  through  sev- 
eral different  acids  liberates  equal  amounts  of  hydrogen  from  each. 
This  is  true,  whether  the  passage  of  the  given  quantity  of  electricity 
is  compressed  into  a  brief  time  in  one  case  and  spread  over  a  longer 
time  in  another,  or  is  uniform  in  all  cases  compared.  It  is  irrespec- 
tive of  the  state  of  dilution  and  of  the  temperature  of  each  acid. 
Thus  two  moles  of  hydrochloric  acid  HC1  are  always  decomposed  for 
every  one  of  sulphuric  acid  H2S04  by  the  same  current.  Similarly,  if 
in  different  cells  we  place  solutions  of  substances  like  sodium  chloride 
NaCl,  cupric  chloride  CuCl2,  antimony  chloride  SbCls,  ferrous  chloride 
FeCl2,  and  ferric  chloride  FeCls,  equal  amounts  of  chlorine  are 
liberated  by  currents  of  equal  strength  in  the  same  time  in  each. 

If  we  consider  the  relation  of  these  facts  to  the  equivalence  of  the 
materials  liberated  in  any  one  cell,  it  will  be  evident  that  when  one 
gram  of  hydrogen  is  liberated  from  each  of  the  two  acids  mentioned 
above,  one  equivalent  of  chlorine  will  be  set  free  in  the  one  cell,  and 
one  equivalent  of  SCX,  or  half  the  weight  represented  by  the  formula, 
will  be  set  free  in  the  other.  Similarly,  with  the  chlorides  of  the  second 
and  third  metals,  while  35.46  g.  of  chlorine  are  being  liberated  in  each 
cell,  the  quantities  of  the  metal  set  free  will  be,  of  copper  one-half  of 
63.57  g.,  and  of  antimony  one-third  of  120.2  g.  Finally,  with  the 
two  iron  salts,  the  quantities  of  iron  liberated  by  the  same  current 
will  be  one-half  and  one-third  of  55.84  grams,  respectively. 

The  simplest  way  in  which  to  insure  the  passage  of  precisely  equal 
amounts  of  electricity  through  all  the  cells  is  to  arrange  them  in 
series.  We  know  that  in  such  circumstances  the  quantity  of  electric- 
ity traversing  any  section  of  the  whole  circuit  must  be  the  same  as 
that  traversing  any  other.  In  a  series  of  cells  containing  substances 
like  the  above,  therefore,  during  the  time  that  1.008  g.  of  hydrogen  is 


IONIZATION 


351 


being  set  free,  we  shall  have  liberation  of  the  equivalent  quantities  of 
each  of  the  other  ions  (Fig.  104) .  Thus  the  second  part  of  Faraday's 
law  states  that:  Equal  quantities  of  electricity  liberate  chemically 
equivalent  quantities  of  the  ions  (equivalent,  p.  63,  not  atomic  or 
molecular). 


__   HC1        H2804       NaCl       CuCl,       SbCl,       FeCl2       Fed 


Cl  H     |80tNa      Cl  JCu     Cl  JSb     Cl  iFe      Cl 


Fia.  104. 

In  this  connection  we  note  that  the  ionizations  of  ferrous  chloride 
and  of  ferric  chloride  are  written  as  follows: 


FeCl2  ^  FC++  +  2C 


FeCl3  ^ 


+  301 


Ferrous-ion  Fe4"4"  and  ferric-ion  Fe4"4"4",  although  differing  only  in  the 
quantity  of  the  charge  carried  by  each  ion,  have  entirely  distinct 
chemical  properties.  In  writing  equations  involving  ions,  care  must 
always  be  taken  to  make  the  numbers  of  +  and  —  charges  equal. 

Application:  The  Explanation  of  Electrolysis.  —  A  battery 
is  a  machine  which  maintains  two  points,  its  poles,  or  two  wires  con- 
nected with  them,  at  a  con-  Cathode  + 

< — cation=Ag 
anion=NO3 — > 


Anode 


O 

o 

O 


e©e 


c 


stant  difference  of  potential. 

One  cell  of  a  lead  storage  bat- 

tery, for  example,  maintains  a 

potential  difference  of  about 

two  volts.     When   the   wires 

are  joined,   directly   or   indi- 

rectly, the  poles  are  immedi- 

ately discharged,  but  the  cell 

continuously    reproduces    the 

difference  in  potential  by  gen- 

erating fresh  electricity.     Now  the  effect  of  immersing  two  plates, 

one  of  which  is  kept  by  the  battery  at  a  definite  positive  potential 

and  the  other  at  a  definite  negative  potential,  into  a  liquid  filled  with 

floating  multitudes  of  minute  bodies,  already  highly  charged,  may 

easily  be  foreseen. 


Fia.  105. 


352  INORGANIC  CHEMISTRY 

The  figure  (Fig.  105)  will  convey  some  idea  of  the  behavior  of  the 
parts  of  a  system  such  as  we  have  imagined.  The  electrodes  are 
marked  —  and  +.  The  negatively  charged  plate  attracts  all  the 
positively  charged  particles  in  the  vessel  and,  although  these  particles 
were  in  continuous  and  irregular  motion,  they  at  once  begin  to  drift 
toward  the  plate  in  question.  On  the  other  hand,  the  negatively 
charged  particles  are  repelled  by  this  plate  and  attracted  by  the 
positive  plate,  so  that  they  drift  in  the  opposite  direction.  Those 
which  are  nearest  each  plate,  on  coming  in  contact  with  it,  will  lose 
their  charges  of  electricity,  turning  thereby  into  the  ordinary  free 
forms  of  the  matter  of  which  they  are  composed.  The  continuous 
removal  of  the  electrical  charges  of  the  plates  through  contact  with 
ions  of  the  opposite  charge  furnishes  occasion  for  recharging  of  the 
plate  from  the  battery,  and  thus  gives  rise  to  a  continuous  current  in 
each  wire.  Again,  the  continuous  drifting  of  positively  and  nega- 
tively charged  particles  in  opposite  directions  through  the  liquid, 
constitutes  what,  in  the  view  of  all  external  means  of  observation, 
appears  to  be  an  electrical  current  in  the  liquid  also.  A  magnetized 
needle,  for  example,  which  is  deflected  when  brought  near  one  of  the 
wires  of  the  battery,  is  influenced  in  the  same  way  by  being  brought 
over  the  liquid  between  the  electrodes.  The  illusion,  so  to  speak,  of 
an  electric  current  is  complete,  although  in  reality  it  is  a  convection  of 
electricity  that  is  taking  place.  Furthermore,  the  quantity  of  elec- 
tricity being  transported  across  any  section  of  the  whole  system  is  the 
same  as  that  across  any  other,  whether  this  section  be  taken  through 
one  of  the  wires,  through  the  electrolyte,  or  even  through  the  battery 
at  any  point.  As  fast  as  the  ions  are  thus  annihilated  as  such,  the 
undissociated  molecules  (mingled  with  the  ions,  but  not  shown  in  the 
figure)  dissociate  and  produce  fresh  ones,  as  in  all  chemical  equilibria. 
Eventually,  by  continuing  the  process  long  enough,  if  the  substances 
set  free  are  actually  deposited  and  do  not  go  into  solution  again  in  any 
form,  the  liquid  can  be  entirely  deprived  of  the  solute  which  it  contains. 

The  analogy  to  the  transportation  of  a  fluid  like  water  is  noticeable, 
although  not  complete.  Water  may  be  transported  in  three  ways. 
It  may  flow  through  a  pipe,  it  may  pass  by  pouring  freely  from  one 
container  to  another,  and  it  may  be  carried  in  vessels.  Thus  a  stream 
of  water,  essentially  continuous,  might  be  arranged,  in  which  part  of 
the  passage  took  place  through  the  pipes,  part  by  pouring  from  the 
pipes  into  buckets,  and  part  by  the  carrying  of  those  buckets  between 
the  ends  of  the  pipes.  The  quantity  of  water  passing  a  given  point  per 
minute  in  this  system  would  be  the  same  at  every  part,  although  the 


IONIZATION  353 

actual  method  by  which  the  water  was  transported  past  the  various 
points  might  be  different.  In  such  a  disjointed  circuit  we  suppose  the 
electricity  to  move  when  carried  from  a  battery  through  an  electrolytic 
cell.  It  flows  in  the  wire,  passes  by  discharge  between  the  pole  and 
the  ion,  and  is  transported  upon  the  ions  in  the  liquid.  The  parallel 
is  imperfect,  however,  because  we  have  used  the  conception  of  two 
electric  fluids  and  because  the  ions  are  already  charged  in  the  solution, 
and  before  any  connection  with  the  battery  i$  made.  They  do  not,  so  to 
speak,  transport  the  electricity  of  the  battery,  but  their  own. 

The  harmony  between  the  quantity  of  electricity  and  the  chemi- 
cal valence  of  the  material  which  it  liberates  is  complete.  The 
picture  which  the  process  of  electrolysis  in  a  series  of  cells  (p.  351) 
presents  to  our  minds  is  very  interesting.  The  progress  of  the  elec- 
tricity through  the  series  is  accompanied  by  a  simultaneous  discharge 
in  all  the  cells  of  chemically  corresponding  numbers  of  ions.  For 
every  atom  of  antimony  that  is  liberated  in  one  cell,  three  atoms  of 
chlorine,  three  atoms  of  hydrogen,  and  one  atom  of  ferric  iron,  are  set 
free  at  the  same  time.  For  two  atoms  of  ferric  iron,  three  atoms  of 
ferrous  iron  and  three  atoms  of  copper  are  deposited.  Even  in  the 
battery  which  generates  the  current,  the  chemical  changes  taking  place 
proceed  atom  for  atom  and  valence  for  valence  in  unison  with  those  in 
the  cells  on  the  circuit.  For  example,  if  the  battery  contains  zinc 
plates,  for  every  atom  of  zinc  that  dissolves,  one  of  copper  and  two  of 
chlorine  will  be  liberated  in  one  of  the  cells.  Our  imaginary  mecha- 
ijsm  thus  puts  all  the  processes  going  on  in  the  circuit  in  the  light  of 
movements  of  the  parts  of  a  perfectly  adjusted  and  interlocked 
machine. 

Questions  Suggested  by  this  Explanation.  —  The  question 
was  raised  (p.  339),  as  to  how  we  can  imagine  separate  atoms  of 
sodium  to  exist  in  water  without  acting  upon  it,  as  the  metal  sodium 
usually  does.  But  the  ions  of  sodium  in  sodium  chloride  solution  are 
not  metallic  sodium.  They  bear  large  charges  of  electricity.  They 
possess  an  entirely  different,  and  in  fact,  by  measurement,  much 
^mailer  amount  of  chemical  energy  than  free  sodium.  And  the 
properties  of  a  substance  are  determined  as  much  by  the  energy  if 
contains  as  by  the  kind  of  matter.  Metallic  sodium  and  ionic  sodium 
are,  simply,  different  substances. 

Besides,  when  metallic  sodium  acts  on  water,  it  turns  into  the 
ionic  sodium  of  sodium  hydroxide  (p.  115):  Na+  +  OH~<=±NaOH. 
Ionic  sodium  Naf  from  sodium  chloride  js,  therefore,  already  in 


354  INORGANIC  CHEMISTRY 

the  very  state  which  metallic  sodium  reaches  by  interaction  with 
water,  and  is  in  no  need  of  trying  to  enter  that  state. 

2.  We  think  of  hydrogen  chloride  and  common  salt  as  exceedingly 
stable  substances,  and  are  averse  to  believing  that  precisely  these  com- 
pounds should  be  highly  dissociated  by  mere  solution  in  water.     But 
it  must  be  remembered  that  in  solution  they  undergo  chemical  change 
very  easily,  and  it  is  only  in  the  dry  form  that  they  show  unusual 
stability. 

3.  Again,  why  do  not  the  ions  combine,  in  response  to  the  attrac- 
tions of  their  charges?    The  answer  is  that  they  do  combine,  but  the 
rate  at  which  combination  takes  place  is  no  greater  than  that  at  which 
the  molecules  decompose,  so  that  on  the  whole  the  proportion  of  ions 
to  molecules  remains  unchanged. 

4.  It  might  appear  that  the  idea  that  bodies  could  retain  high 
charges  in  the  midst  of  water  is  contrary  to  all  experience.     It  must 
be  remembered,  however,  that  the  molecular,  pure  water,  which 
separates  the  ions  from  one  another,  is  a  perfect  nonconductor.     The 
moisture  which  covers  electrical  apparatus  and  causes  leakage  of 
static  electricity  is  not  pure  water,  but  a  dilute  solution  containing 
carbonic  acid  (p.  197)  and  materials  from  the  glass  of  which  the 
apparatus  is  made  (p.  143) .     It  conducts  away  the  charge  electro- 
lytically,  by  means  of  the  ions  it  contains,  and  not  by  itself  acting  as 
a  conductor. 

5.  Finally,  when  we  dissolve  an  electrically  neutral  salt  in  water, 
whence  do  the  radicals  obtain  the  electric  charges?    We  now  know 
that  an  atom,  say  of  sodium,  contains  a  minute  nucleus  of  positive 
electricity,  which  contains  most  of  the  mass  of  the  atom.     Outside 
of  this  nucleus,  there  are  particles  of  negative  electricity,  called 
electrons  (q.v.),  each  having  a  mass  about  one-seventeen  hundredth 
(nW)  of  that  of  an  atom  of  hydrogen.     An  ion  of  chlorine  (Cl~) 
consists,  therefore,  of  an  atom  of  chlorine  plus  one  electron  (Cl  -f-  e). 
An  ion  of  sodium  is  an  atom  of  sodium  minus  one  electron  (Na  —  e) 
and  has  thus  an  excess  of  one  unit  positive  charge  in  the  nucleus. 
When  these  two  ions  combine,  the  resulting  molecule  NaCl  is  neutral. 

According  to  this  view,  a  battery  is  a  source  of  electrons,  which 
pass  in  a  stream  along  the  negative  wire.  When  the  electrons  reach 
the  electrode  in  an  electrolytic  cell,  they  (being  negative)  attract  the 
positively  charged  ions  and,  when  the  latter  touch  the  electrode,  the 
missing  one  or  more  electrons  are  supplied  to  each  ion,  and  the 
material  of  the  ion  is  then  neutral  and  free: 

2H+  +  2e->H2T,     or    Cu++  +  2e  -»  Cu  1  . 


IONIZATION  355 

The  negative  ions  are  repelled  by  the  negative  electrode  .and  move 
toward  the  positive  one.  Also,  if  several  positive  ions  could  be 
rendered  neutral  and  free,  there  would  be  an  excess  of  negative 
charges  in  the  solution.  To  avoid  this,  the  negative  ions  in  contact 
with  the  positive  electrode  give  up  an  equivalent  number  of  electrons 
to  the  plate  and  wire :  2C1~  — »  Ck  T  +  2e.  Thus,  the  circulation 
of  the  electrons  is  completed,  and  the  current  of  electrons  flows  in  all 
parts  of  the  system. 

Resume  and  Nomenclature.  —  The  dissociation  of  molecules 
into  ions  is  named  ionization.  The  substances  of  the  three  classes 
which  alone  are  ionized  are  designated  ionogens.  An  ion  may  be 
defined  as  a  molecule  bearing  negative  or  positive  charges  of  elec- 
tricity in  proportion  to  its  valence,  ;vad  formed  through  the  dissocia- 
tion of  an  ionogen  by  a  solvent  like  water. 

The  solution  of  an  ionized  substance  is  called  an  electrolyte,  and  often  this 
term  is  applied  also  to  acids,  bases,  and  salts  themselves,  because,  when  dissolved, 
they  produce  electrolytes.  This  is  rather  a  confusing  metonymy,  however,  be- 
cause these  bodies  by  themselves  are  not  conductors.  This  use  of  the  term  also 
introduces  obscurity  because  it  connects  the  ionization  with  electrolysis  and  al- 
ways conveys  the  impression  that  the  latter  produces  the  former.  The  electro- 
lytic property -of  ions  is  only  one  amongst  many  special  properties  of  electrolytes, 
and  the  majority  of  these  properties  are  chemical  and  have  nothing  to  do  with 
electrolysis.  Hence  we  have  preferred  the  more  general  word  "ionogen." 

Each  molecule  of  the  solute  gives  two  kinds  of  ions  with  opposite 
charges.  These  two  are  forthwith  distinct  and  independent  sub- 
stances, save  that  the  attractions  of  the  charges  prevent  any  con- 
siderable separation  by  diffusion.  They  differ  from  non-ionic. sub- 
stances of  the  same  material  composition  when  such  are  known. 
The  electrical  charge  is  one  of  the  essential  constituents  and,  when 
it  is  removed,  the  properties  alter  entirely.  Thus  we  have  two 
kinds  of  hydrogen  —  gaseous,  molecular  hydrogen  (H2),  and  ionic 
hydrogen  (H+)  —  with  entirely  different  chemical  properties  (p.  340). 

The  radicals  and  their  chemical  behavior  are  real,  and  all  the 
peculiarities  of  aqueous  solutions  of  acids,  bases,  and  salts  are  ex- 
perimental facts.  We  now  have  experimental  knowledge  of  the 
minute  parts  of  bodies.  Molecules  are  units  which  are  not  commonly 
disintegrated  by  vaporization  (p.  233);  ions,  those  which  are  not 
commonly  disintegrated  in  double  decomposition  in  solution;  atoms, 
those  which  are  not  commonly  disintegratedjki  any  chemical  action. 
But  there  are  exceptions  in  each  of  the  three  cases.  The  ionic 


356 


INORGANIC  CHEMISTRY 


explanation  was  first  suggested  as  an  hypothesis  by  Svante  Arrhenius, 
a  Swedish  chemist,  in  1887.  From  the  appearance  of  his  remarkable 
memoir  we  date  the  great  development  which  the  study  of  solutions* 
has  undergone  in  recent  years. 

It  is  worth  noting  that  the  quantities  expressed  by  the  formulae  Al,  Ca,  and  K, 
when  existing  as  ions,  produce  equal  osmotic  pressures,  and  have  equal  effects 
upon  the  freezing-  and  boiling-points.  This  is  a  further  justification  for  our 
choice  of  chemical  unit  quantities  of  the  elements  (atomic  weights),  for  the 
atomic  weights  have  these  properties  in  common,  and  equivalents,  of  course,  do 
not  (cf.  p.  245). 

Since  ionic  hydrogen,  ionic  chlorine,  etc.,  are  entirely  different  in 
physical  and  chemical  properties  from  the  corresponding  free  ele- 
ments, they  should  receive  separate  names.  When  it  is  incon- 
venient to  say  "ionic  hydrogen,"  " ionic  nitrate  radical"  (NOs""), 
etc.,  the  following  names  will  be  used  for  the  ionic  substances: 


Symbol. 

Name  of 
Substance. 

Anion  of 

Symbol. 

Name  of 
Substance. 

Cation  of  Salts  of 

S04= 

Sulphate-ion 

Sulphates 

Na+ 

Sodium-ion 

Sodium 

cr 

Chloride-ion 

Chlorides 

Fe+++ 

Ferric-ion 

Ferric  iron 

Fso4~ 

Hydrosulphate-ion 

Bisulphates 

NH4+ 

Ammonium-ion 

Ammonium 

OK 

Hydroxide-ion 

Hydroxides 

Fe++ 

Ferrous-ion 

Ferrous  iron 

(bases) 

H+ 

Hydrogen-ion 

Hydrogen  (acids) 

In  using  these  terms,  note  that  sodium-ion  (with  the  hyphen)  is  the 
name  of  the  substance,  and  not  of  the  charged  atom.  When  speaking 
in  terms  of  ions  as  particles,  therefore,  we  may  not  say  "a  sodium- 
ion,"  any  more  than  we  should  say  "an  ionic  sodium"  or  "ionic 
sodiums."  To  describe  the  charged  molecule,  we  must  write  "a 
sodium  ion,"  "sodium  ions,"  "chlorate  ions,"  etc. 

Faraday  distinguished  by  name  the  two  kinds  of  material  which 
proceed  with  and  against  the  positive  current.  His  terminology  is 
still  used.  Ions  which  proceed  in  the  same  direction  as  the  positive 
current  (Fig.  105,  p.  351)  are  called  cations  (Gk.  /card,  down).  Such 
are  H+,  GU++,  K+,  NH4+.  They  are  metallic  elements,  or  groups 
which  play  the  part  of  a  metal.  The  electrode  (Gk.  SSos,  a  path) 
upon  which  they  are  deposited,  the  negative  electrode,  is  spoken  of  as 
the  cathode  (Gk.  3  KaOoSos,  the  way  down). 


*  The  Scientific  Memoirs,  No.  IV  (American  Book  Company),  is  a  reprint 
of  the  fundamental  papers  by  Raoult,  van't  Hoff,  and  Arrhenius. 


IONIZATION  357 

The  particles  which  move  in  the  direction  of  the  negative  current, 
and  against  that  of  the  positive,  are  named  anions  (Gk.  avd,  up).  The 
ions  Cl~,  NOs",  SC>4=,  Mn04~  are  of  this  kind.  They  are  usually  com- 
posed of  non-metals,  although  sometimes,  as  in  MnCU",  the  constitu- 
ents may  be  partially  metallic.  They  are  set  free  at  the  positive  elec- 
trode, which  is  therefore  named  the  anode  (Gk.  ^  avoSos,  the  way  up). 
Chemists  speak  of  metallic  and  non-metallic  elements  as  positive 
and  negative  elements,  respectively  (cf.  p.  150),  even  when  electrical 
relations  are  not  directly  in  question,  and  ions  are  not  concerned. 

In  order  that  this  idea  may  be  carried  out  consistently,  the  libera- 
tion of  any  of  these  ionic  materials  at  one  electrode  in  electrolysis  is 
written  as  follows: 

Ag*  +  9  ->  Ag,  2CP  +  2©  -»  C12. 

Here  0  and  ©  represent  the  unit  quantities  of  negative  and  posi- 
tive electricity  furnished  by  the  battery  to  the  electrodes  and  de- 
stroyed by  opposite  charges  upon  the  ions. 

The  solvents  which  produce  ionization  are  the  associated  liquids 
(p.  282),  like  water  (p.  202),  liquefied  ammonia,  and  alcohol.  Those 
which  do  not  produce  ionization  are  the  non -associated  liquids,  like 
benzene,  toluene,  and  carbon  disulphide.  This  likewise  is  explained 
by  the  fact  that  the  former  have,  by  measurement,  high  dielectric 
constants,  and  so  are  able  to  hold  charged  bodies  apart,  while  the 
latter  have  low  ones.  • 

Actual  Quantities  of  Electricity  Concerned.  —  The  units 
of  electrical  energy  are  the  coulomb,  which  is  the  unit  of  quantity, 
and  the  volt,  which  is  the  unit  of  difference  of  potential  (or  pressure, 
so  to  speak).  Faraday's  law  has  to  do  only  with  the  former.  Equal 
numbers  of  coulombs  liberate  equivalent  weights  of  all  elements, 
but  different  voltages  are  required  to  decompose  different  compounds, 
according  to  their  stability  (see  Chap.  XXXVIII). 

To  liberate  1.008  g.  of  hydrogen,  or  one  equivalent  of  any  other 
element,  96,504  coulombs  of  electricity  are  needed.  The  charges 
on  1.008  g.  of  hydrogen  ions  must,  therefore,  equal  this  amount. 
There  are  6.07  X  1023  molecules  of  hydrogen  in  22.4  liters  (H2)  and 
therefore  in  2.016  g.  of  the  gas.  A  simple  calculation  shows  there- 
fore that  each  coulomb  is  distributed  over  about  63  X  1017  ions  of 
hydrogen. 

A  current  of  1  coulomb  per  second  is  called  1  ampere.  Thus, 
the  current  passing  through  a  1-amp.  lamp  (or  2  half-ampere,  50 


358  INORGANIC  CHEMISTRY 

watt*  lamps  in  parallel)  will  liberate  1.008  g.  (11.2  liters)  of  hydro- 
gen in  96,504  seconds,  or  26  hours  and  49  minutes.  The  same 
current  will  liberate  107.88  g.  of  silver  (Ag1),  or  31.78  g.  of  copper 
(Cun/2)  from  a  cupric  salt  in  the  same  time.  A  current  of  5 
amperes  will  accomplish  the  same  result  in  one-fifth  of  the  time. 
In  consequence  of  this,  the  liberation  of  hydrogen  from  a  dilute 
acid  by  electrolysis,  or  the  deposition  of  silver  or  copper,  is  used  as 
a  means  of  measuring  quantities  of  electricity.  The  volume  of 
hydrogen  collected,  or  the  increase  in  weight  of  the  negative  electrode 
in  a  cell,  called  under  such  circumstances  a  voltameter  (measurer  of 
coulombs,  not  volts),  is  a  measure  of  the  quantity  of  electricity  which 
passes  around  the  whole  circuit  of  which  it  forms  a  part. 

Applications:  Ionic  Equilibrium.  —  Since  the  ions  are  chemi- 
cally different  from  their  parent  molecules,  their  formation  repre- 
sents a  variety  of  chemical  change.  The  change  may  not  involve 
any  chemical  interaction  with  the  water.  It  is  simply  a  dissocia- 
tion, i.e.,  reversible  decomposition  of  the  dissolved  substance. 

From  the  fact  that  the  proportion  of  molecules  ionized  is  shown 
to  become  greater  as  more  and  more  of  the  solvent  is  added  (p.  337), 
and  that  removal  of  the  solvent  diminishes  the  proportion  of  ions  to 
molecules,  and  finally  leaves  the  substance  entirely  restored  to  the 
molecular  condition,  we  know  that  this  is  a  reversible  action  and 
therefore  a  true  dissociation.  The  molecules  and  their  ions  adjust 
themselves  like  the  components  in  any  case  of  chemical  equilibrium 
(pp.  289-307): 


These  equilibria  are  all  of  precisely  the  same  nature  as  that  of 
phosphorus  pentachloride  vapor  (p.  299),  and  the  discussion  of  the 
latter  should  be  reexamined  and  applied  by  the  reader.  The  sole 
difference  is  that  here  change  in  volume  is  effected,  not  by  compression 
or  by  release  of  pressure,  but  by  removing  or  adding  water.  The 
adjustment  to  a  condition  of  equilibrium,  however,  seems  to  be  instan- 
taneous where  ions  are  concerned,  while  in  other  chemical  actions  it 
always  takes  a  perceptible,  and  often  a  considerable  interval  of  time. 
The  chemical  behavior  of  substances  in  ionic  equilibrium  will  be 
discussed  in  the  next  chapter  (see  p.  377). 

*  No.  amperes  X  no.  volts  =  no.  watts.  Hence,  with  100  volts,  a  50-watt 
lamp  carries  one-half  ampere. 


IONIZATION  359 

*  The  mode  of  formulation  previously  used  (p.  298)  may  be 
employed  here.  If  [NaCl],  [Na+],  and  [Cl~]  stand  for  the  molec- 
ular concentrations  (numbers  of  moles  per  liter)  at  equilibrium  of 
the  molecules,  and  the  two  ions,  respectively,  we  have  an  equilib- 
rium constant  (cf.  p.  298),  in  this  case  called  the  ionization  constant: 

[Na+]  X  [C1-] 


= 


[NaCl] 


When  we  dissolve  a  single  substance  which  gives  only  two  ions,  the 
molecular  concentrations  of  the  ions  are  necessarily  equal.  When 
some  other  ionogen  with  a  common  ion  is  present,  however,  the 
values  of  [Na+]  and  [Cl~]  will  be  different.  • 

The  Electrical  Energy  Required  to  Decompose  Different 
Compounds.  —  Chemical  compounds  are  of  very  different  degrees 
of  stability,  and  hence  the  quantities  of  energy,  electrical  or  otherwise, 
required  to  decompose  them  vary  widely.  Thus,  hydrogen  chloride 
is  very  stable,  while  hydrogen  iodide  is  easily  decomposed  by  heating. 
The  disunion  of  one  mole  (equivalent  quantities)  of  these  substances 
in  aqueous  solution  absorbs  39,300  cal.  and  13,100  cal.  of  heat  energy, 
respectively.  Hence,  although  equal  quantities  of  electricity  (96,504 
coulombs  in  each  case)  perform  this  office,  very  unequal  amounts  of 
electrical  energy  are  used  up  in  the  electrolysis. 

The  energy  in  a  stream  of  water  is  represented  by  the  product  of 
the  quantity  passing  a  given  section  and  the  pressure  or  head  of  water 
available  at  that  point.  If  the  pressure  is  low,  the  work  that  can  be 
done  will  be  small,  even  if  the  quantity  flowing  is  great.  So  electrical 
energy  is  expressed  by  the  product  of  the  current  strength,  or  quantity 
of  electricity  passing  per  second  during  a  certain  period  of  time,  and 
the  electromotive  force.  The  latter  corresponds  to  pressure,  and  is 
denned  by  the  difference  in  potential  of  two  points  in  the  circuit 
between  which  the  energy  is  being  used  up. 

Now,  in  the  series  of  cells  which  was  described  (p.  351),  each  cell, 
while  being  traversed  by  the  same  quantity  of  electricity  as  any  of 
the  others,  cuts  down  the  electromotive  force  of  the  current  in  propor- 
tion to  the  amount  of  energy  consumed  by  the  decomposition  going  on 
within  it.  Hence,  while  a  voltmeter  will  show  no  difference  in  poten- 
tial between  two  neighboring  parts  of  the  heavy  wires  used  as  connec- 

*  The  content  of  this  paragraph  is  referred  to  in  Chap.  XX,  but  is  not  em- 
ployed systematically  until  Chap.  XXXIV  is  reached. 


360  INORGANIC  CHEMISTRY 

tions,  for  no  work  is  being  done  in  the  wires,  it  will  show  a  considerable 
difference  in  potential  between  two  points  which  are  separated  by  one 
of  the  cells. 

A  system  of  cars  hauled  by  a  cable  is  analogous  to  our  set  of  cells 
and  more  familiar.  When  clutched  to  the  cable,  all  the  cars  move 
with  equal  speed,  but,  being  loaded  with  different  numbers  of  passen- 
gers, take  very  different  amounts  of  power  from  the  moving  cable. 

We  should  infer  from  this,  that  to  decompose  every  electrolyte,  a 
current  of  a  certain  minimum  electromotive  force,  sufficient  to  furnish 
the  fall  in  potential  necessitated  by  the  chemical  change,  which  would 
be  different  in  different  cases,  would  be  required.  This  is  found  to  be 
the  case.  For  the  easy  decomposition  of  sulphuric  acid  and  liberation 
of  the  products  an  electromotive  force  of  at  least  1.92  volts  is  necessary, 
for  hydrochloric  acid  1.41  volts,  for  hydriodic  acid  0.53  volts,  for  zinc 
sulphate  2.7  volts,  and  for  silver  nitrate  0.70  volts.  When  we  use  a 
current  of  electromotive  force  falling  short  of  that  specified,  we  find 
that  the  flow  of  electricity  is  interrupted.  The  electrolytic  cell  practi- 
cally forms  a  break  in  the  circuit  (see  Chap.  XXXVIII). 

Polarization.  —  It  is  found  that  when  plates  of  platinum,  a 
metal  which  is  not  acted  upon  by  the  liberated  radicals,  are  used,  the 
products  of  electrolysis  accumulate  on  the  electrodes  and  tend  to 
produce  a  reverse  current  (see  Electromotive  chemistry) .  The  cell  is 
said  to  be  polarized.  Thus,  after  hydrochloric  acid  has  been  elec- 
trolyzed  for  a  few  moments,  hydrogen  and  chlorine  adhering  to  the 
two  platinum  plates  set  up  this  current.*  If  the  battery  is  dis- 
connected, the  electrolytic  cell  becomes  for  a  brief  time  itself  a  battery, 
the  re-ionization  of  the  hydrogen  and  chlorine  (reproducing  hydro- 
chloric acid)  furnishing  the  energy.  It  is  the  continuous  overcoming 
of  this  reverse  current,  and  prevention  of  the  re-ionization,  that 
demands  the  minimum  electromotive  force  (here  1.41  volts)  of  which 
mention  has  just  been  made. 

It  is  possible  to  arrange  cells  in  which  no  polarization  can  take 
place.  Thus,  when  we  electrolyze  cupric  sulphate  between  copper 
electrodes,  the  copper  is  deposited  upon  one  plate  and  the  SO4  removes 
the  copper  from  the  other  plate,  forming  cupric  sulphate,  thus  restor- 
ing the  electrolyte  to  its  original  condition.  The  only  difference  is 
that  a  portion  of  the  copper  has  been  deposited  on  one  pole  and  an 
equivalent  amount  has  been  removed  from  the  other  (see  Copper 

*  If  copper  plates  are  used,  cupric  chloride  is  formed  at  the  positive  plate 
(anode),  and  no  polarization  can  occur  at  that  plate. 


IONIZAT1ON  361 

refining).  With  such  cells,  no  minimum  difference  in  potential  is 
required  to  effect  electrolysis,  for  there  is  no  polarization  current  to  be 
overcome.  The  feeblest  electric  current  will  produce  continuous,  if 
slow,  chemical  change. 

This  result  is  extremely  interesting,  for  it  shows  that  the  operation 
of  migration  in  itself  does  not  require  the  consumption  of  much 
energy.  If  the  molecules  were  actually  torn  apart  by  the  electricity, 
then  all  electrolytic  operations  would  require  a  minimum  electro- 
motive force  for  their  maintenance.  The  fact  just  stated,  therefore, 
is  significant,  for  it  confirms  the  present  views  in  regard  to  the  theory 
of  solutions.  This  fact  is  in  agreement  with  the  belief  that  the  actual 
production  of  the  ions  is  accomplished  by  the  water  in  advance,  and 
quite  independently  of  the  use  of  electricity,  and  that  the  sole  function 
of  the  electricity  in  the  process  of  electrolysis  within  the  solution  con- 
sists in  the  pilotage  of  the  ions  in  reverse  directions  according  to  their 
charges,  an  operation  which  necessarily  consumes  but-  little  energy. 
The  friction  alone  of  the  moving  ions  has  to  be  overcome.  It  makes 
clear  the  fact  that  it  is  only  when  the  chemical  change  in  the  cell 
involves  the  actual  decomposition  of  some  material,  accompanied  (as 
in  the  electrolysis  of  hydrochloric  acid)  by  the  final  delivery  of  the 
constituents  in  the  free  state,  that  considerable  consumption  of 
electrical  energy,  proportional  to  the  extent  of  the  chemical  change, 
must  take  place. 

Applications:  The  Conductivity  of  Electrolytes  for  Elec- 
tricity. —  The  facility  with  which  equi-molar  solutions  of  different 
substances  conduct  electricity,  when  they  are  placed  under  like  condi- 
tions, depends  jointly,  (l)  on  the  degree  of  ionization,  (2)  on  the  speed 
with  which  the  ions  move,  and  (3)  on  the  valence  of  the  ions. 

This  is  easily  explained.  Equal  numbers  of  the  original  molecules 
are  used  in  making  the  different  solutions  being  compared  (equi-molar 
solutions).  Evidently  the  greater  the  percentage  of  the  molecules  of 
the  solute  which  is  ionized  (1),  the  larger  the  number  of  the  ions 
present,  and  the  greater  the  facilities  for  transporting  electricity 
through  the  solution.  Also,  with  equal  numbers  of  ions,  those  which 
move  fastest  (2)  will  carry  the  most  electricity  through  the  solution. 
Finally,  with  equal  numbers  of  ions,  and  equal  speeds,  ions  with  higher 
valence  (3),  and  therefore  more  unit  charges  per  ion,  will  carry  more 
electricity  than  will  ions  with  only  one  charge  per  ion  (univalent 
ions) .  The  case  is  like  ferrying  an  army  across  a  river.  The  number 
of  soldiers  passing  per  hour  will  depend,  (1)  upon  the  number  of  boats 


362  INORGANIC  CHEMISTRY 

available,  (2)  upon  the  speed  with  which  each  boat  moves,  and  (3) 
upon  the  capacity  of  each  boat. 

When  equivalent  instead  of  equi-molar  amounts  are  compared, 
the  last  of  these  factors,  the  valence,  drops  out  of  consideration. 

The  most  highly  dissociated  acids,  as  we  should  expect,  since  they 
give  large  numbers  of  the  speedy  hydrogen  ions,  are  the  best  con- 
ductors. The  highly  ionized  bases,  such  as  potassium  and  sodium 
hydroxides,  come  next.  The  best  conductors  among  salts  fall  con- 
siderably behind  both  of  these,  because,  although  their  degrees  of 
ionization  may  not  be  less  than  those  of  the  best  conducting  acids  and 
bases,  their  ions  all  move  more  slowly  than  do  hydrogen-ion  and 
hydroxide-ion.  On  the  other  hand,  concentrated  solutions  all 
conduct  badly,  relatively  to  the  number  of  molecules  originally  used 
in  making  them,  because  only  that  proportion  of  the  substance  which 
is  ionized  contributes  to  the  conduction  (p.  358).  All  this  is  just  what 
we  should  expect,  in  view  of  our  hypothesis,  for  the  passage  of  the 
electricity  must  be  dependent  upon  the  frequency  with  which  dis- 
charges of  the  ions  upon  the  electrodes  occur,  and  this,  in  tui  i,  must 
depend  upon  both  the  concentration  and  the  speed  of  the  ions.  To 
returh  to  an  analogy  used  once  before,  the  rate  at  which  a  fluid  can  be 
transferred  between  two  reservoirs  must  depend  upon  the  denseness 
of  the  array  of  buckets  available,  on  the  speed  with  which  they  are 
moved,  and  on  their  individual  capacity. 

Ordinarily,  it  is  the  resistance  which  a  substance  presents  to  the 
passage  of  the  electric  current  which  is  measured.  Obviously,  how- 
ever, for  the  present  purpose  it  is  more  convenient  to  give  expression 
to  the  reciprocal  of  this  value,  which  we  term  the  conductivity.  In 
order  that  the  results  may  have  chemical  significance,  we  express  them 
in  terms  of  the  conducting  power  of  one  gram-equivalent  of  the  com- 
pound dissolved  in  water  and  placed  in  a  narrow  cell  whose  opposite 
walls,  of  great  area  and  situated  one  centimeter  apart,  form  the 
electrodes.  Since  the  water  is  a  nonconductor,  the  conducting  power 
of  the  solution  intervening  between  the  plates  is  a  measure  of  the 
capacity  of  the  dissolved  substance  for  facilitating  the  discharge 
between  the  poles. 

Applications:  To  the  Interpretation  of  Conductivity 
Measurements.  —  We  have  seen  that  when  the  solution  of  an 
ionogen  is  diluted,  the  proportion  of  ions  to  undissociated  molecules 
increases,  while  removal  of  a  part  of  the  solvent  has  the  opposite 
effect  (p.  358).  Now,  a  change  in  the  number  of  ions  naturally 


IONIZATION 


363 


modifies  the  capacity  of  the  liquid  for  carrying  electricity,  so  that 
observation  of  the  changes  in  the  conductivity  of  a  solution,  when 
the  concentration  is  altered,  supplies  the  simplest  means  of  studying 
the  phenomena  of  ionization. 

A  glass  trough  and  amperemeter  *  (Fig.  106)  may  be  used  to  illus- 
trate this  principle.  The  electrodes  are  long  strips  of  copper  foil, 
which  pass  down  at  the  ends  of  the  trough  and  are  situated,  not  one 
centimeter,  but  ten  or  fifteen  centimeters  apart,  in  order  that  the 


Fia  106. 

contents  of  the  vessel  may  be  more  easily  seen.  After  placing  the  two 
instruments  in  circuit  with  a  source  of  electricity,  we  first  pour  very 
pure  water  into  the  cell.  With  this  arrangement,  the  ampere- 
meter does  not  indicate  the  passage  of  any  current  of  electricity. 
Concentrated  (36  per  cent)  hydrochloric  acid  is  next  cautiously 
added  through  a  long-stemmed  dropping  funnel,  so  that  it  forms  a 
shallow  layer  below  the  water.  The  situation  at  this  stage  is  that  a 
definite  amount  of  hydrogen  chloride  dissolved  in  a  small  amount  of 
water  fills  what  was  before  a  nonconducting  gap  in  the  electric  circuit. 
The  deflection  of  the  needle  in  the  amperemeter  indicates  that  a 
certain  current  of  electricity  is  able  to  pass  through  this  acid.  When 
we  now  stir  the  surface  of  the  acid  very  gently  with  a  thin  glass  rod, 
the  amperemeter  instantly  responds,  showing  an  increase  in  conduc- 
tivity. As  we  stir,  the  conductivity  increases,  and  the  increase  ceases 
only  when  the  liquid  has  become  homogeneous.  Introduction  of  an 

*  For  these  experiments  an  amperemeter  of  low  resistance,  0.5-1  ohm,  must 
be  used,  and  a  battery  of  one  or  two  accumulator  cells  is  sufficient. 


364  INORGANIC  CHEMISTRY 

additional  supply  of  water  will  improve  the  conductivity  still  more, 
but  the  effect  becomes  less  and  less,  until  no  change  on  further 
dilution  is  perceptible.  Reasoning  about  these  effects,  we  perceive 
that  the  amount  of  hydrochloric  acid  has  not  altered  during  the 
experiment.  Yet  the  quantity  of  conducting  material  between  the 
electrodes  must  have  become  greater,  for  the  carrying  power  of 
the  whole  has  improved.  We  were  therefore  observing  the  progress 
of  a  chemical  change  of  the  nonconducting  hydrogen  chloride  into 
conducting  materials.  Hydrogen  chloride  molecules  do  not  carry 
electricity  (p.  211),  but  the  hydrogen  and  the  chloride  ions,  into 
which  they  were  gradually  altered  by  chemical  change  during  the  stir- 
ring, do  carry  electricity.  Furthermore,  the  change  practically  ceased 
at  great  dilution,  for  the  dissociation  into  ions  was  then  practically 
complete.  If  we  could  conveniently  have  started  with  only  liquefied, 
dry  hydrogen  chloride  in  the  cell,  we  should  have  observed  the  whole 
range  of  changes  from  zero  to  the  maximum. 

When  a  saturated  solution  of  cupric  chloride  is  used  instead  of  the 
hydrochloric  acid,  dilution  is  accompanied  by  a  similar  improvement 
in  conductivity.  Here  we  notice,  besides,  that  the  yellowish-green 
liquid  with  which  we  start  changes  to  a  pale  blue,  as  the  molecules  of 
cupric  chloride  are  dissociated  and  the  color  of  the  solution  becomes 
more  exclusively  that  of  the  copper  ions.  When  the  solution  has 
become  perfectly  blue,  further  dilution  is  seen  to  affect  the  conduc- 
tivity but  slightly. 

The  approach  to  a  maximum  of  conductivity  reached  in  these  two 
cases  indicates  that  practically  the  whole  of  the  material  has  assumed 
the  ionic  form.  Theoretically  the  absolute  maximum  would  be 
reached  at  infinite  dilution.  The  conductivity  of  the  same  amount  of 
substance  in  more  limited  dilution  is  that  of  the  proportion  of  ions 
corresponding  to  this  dilution,  since  the  complete  molecules,  still 
present,  are  without  influence  on  conductivity.  Thus  the  ratio  of  the 
conductivity  at  a  given  dilution  to  the  maximum  conductivity  is  equal 
to  the  proportion  of  the  whole  material  ionized  at  the  given  dilution. 
From  a  series  of  measurements  for  a  fixed  amount  of  a  substance  at 
different  dilutions,  after  the  results  have  been  plotted,  we  can  usually 
(see,  however,  below)  ascertain  the  limiting,  maximum  conductivity 
by  graphic  extrapolation.  If  X*  is  the  conductivity  of  an  equivalent  of 
the  substance  dissolved  in  v  liters  of  water,  and  X^  the  conductivity 
of  the  same  amount  at  infinite  dilution,  then  \v/\n  is  the  proportion  of 
molecules  completely  ionized  in  the  former  solution.  X»  is  called 
the  equivalent  conductivity  at  the  dilution  v. 


IONIZATION 


365 


The  following  numbers  show  the  equivalent  conductivities  at  18° 
of  solutions  of  four  different  substances,  expressed  in  the  units  always 
employed  for  the  purpose  (which  are  reciprocal  ohms) .  The  symbols 
AO.I,  meaning  1  equivalent  in  0.1  1.;  Xi,  meaning  1  equivalent  in  11.; 
and  so  forth,  denote  the  concentrations. 


Xo.i 

Xi 

KM 

Xioo 

Xoo 
(Calc.) 

Hydrochloric  acid      .    .    . 
Sodium  chloride      .... 

64.4 

301.0 

74.4 

351.0 
92.5 

370.0 
103.0 

384 
110 

Sodium  acetate   

41.2 

61.1 

70.2 

78 

Acetic  acid 

0  05 

1  32 

4  6 

14  3 

(352) 

It  will  be  seen  from  inspection  of  these  figures  that  in  the  case  of 
the  first  three  substances,  the  conductivity  does  not  improve  much 
when  a  solution  containing  one  equivalent  in  10 1.  is  diluted  ten  times, 
and  that  further  dilutions,  no  matter  how  extensive,  produce  a  still 
smaller  effect.  The  case  of  acetic  acid  requires  separate  consideration 
(see  next  section). 

Thus,  measurements  of  conductivity  enable  us  to  study  the  ionic 
decomposition  of  all  ionogens,  and  to  state  accurately  the  fraction 
ionized,  at  each  concentration,  in  solutions  of  every  ionogen.  This 
information  is  obviously  most  valuable,  for  it  places  us  in  a  position 
to  know  the  exact  constitution  of  every  solution  we  use  in  the  labora- 
tory. In  the  next  section  but  one  the  data  on  which  such  knowledge 
can  be  based  is  given.  In  the  following  chapter  the  mode  of  applying 
the  data  is  explained. 

The  Case  of  a  Little- Ionized  Substance.  —  Acetic  acid,  a 
feebly  ionized  substance,  as  the  table  shows,  conducts  very  badly  in 
concentrated  solution  and,  while  the  conductivity  improves  with  dilu- 
tion, it  is  not  possible  experimentally  to  observe  any  approach  to  the 
maximum  conductivity.  The  conductivity,  in  cases  like  this,  is  still 
far  removed  from  the  maximum  at  dilutions  at  which,  with  other  sub- 
stances, the  maximum  is  nearly  attained. 

In  cases  like  that  of  acetic  acid,  the  conductivity  at  infinite  dilu- 
tion cannot  be  estimated  by  extrapolation.  But  fortunately  another 
method  is  available.  The  values  384,  110,  and  78  for  X^  in  the  cases 
of  the  first  three  substances  can  be  reached  by  extrapolation,  and  rep- 
resent the  conducting  powers  of  equal  numbers  of  ions,  for  there  are 
equal  numbers  of  equivalents  present  and  no  molecules  remain  un-ion* 


366  INORGANIC  CHEMISTRY 

ized.  These  values  are  unequal  solely  because  of  the  differing  speeds 
of  the  ions  concerned.  Each  of  them  derives  its  value  from  numbers 
representing  the  relative  speeds  of  the  two  ions  present,  and  must  be 
the  sum  of  these  two  numbers.  If,  therefore,  we  measure  the  relative 
speeds  of  the  two  ions  (p.  347)  of  one  of  the  substances,  we  can  divide 
the  value  of  Xw  in  this  proportion  and  learn  the  part  which  each  ion 
contributes  to  the  total.  Dividing  384  in  this  way  we  get  the  speed 
of  H+  =  318  and  of  Cl~  =  65.9  already  given  (p.  347).  Dividing  XOT 
for  sodium  acetate  (78),  similarly,  we  get  the  speeds  Na+  =  44.4  and 
C02CH3~  =  33.7.  The  speeds  of  CT  and  Na+  together  (65.9  + 
44.4  =  110.3)  must  then  equal  Xw  for  NaCl  and,  as  we  see,  they  do. 
Similarly,  the  speeds  of  H+  and  C02CH3~  together  (318  +  33.7  = 
351.7)  must  equal  XOT  for  HCC^CHs,  although  we  cannot  observe  the 
latter  directly.  This  method  can  be  applied  to  all  of  the  less  highly 
ionized  acids  and  bases,  for  their  sodium  and  potassium  salts  belong 
invariably  to  the  class  of  substances  which  are  most  ionized,  and  for 
which,  therefore,  X^  can  be  determined  accurately  by  extrapolation. 

Composition  of  Solutions  of  lonogens:    Fractions  Ionised. 

—  The  rule,  degree  of  ionization  =  X^/X^  (p.  364),  enables  us  to  cal- 
culate the  value  for  any  dilution,  when  the  necessary  data  are  given. 
We  need  only  the  values  of  the  conductivity  (X»)  for  different  dilutions 
(p.  365)  and  those  of  the  relative  speeds  of  each  kind  of  ions  expressed 
in  the  same  units.  The  latter,  when  added,  give  X^ . 

Thus,  hydrogen  chloride  in  a  solution  containing  1  equivalent  in 
0.1  1.  (365  g.  per  liter),  which  would  be  a  rather  concentrated  hydro- 
chloric acid,  shows  the  fraction  ionized  ||j,  or  0.168  (=  16.8  per 
cent).  Normal  hydrochloric  acid  is  ionized  to  the  extent  of  |^,  or 
O.Z84;  normal  sodium  chloride,  -j^-,  or  0.676;  normal  acetic  acid, 
|f,  0.004  (=  0.4  per  cent). 

The  dilute  acids  used  in  the  laboratory  are  generally  of  six  times 
normal  (QN)  concentration.  But,  often,  we  add  only  a  drop  or 
two  to  a  large  bulk  of  liquid,  so  that  the  acids  are  commonly  very 
dilute  as  actually  employed.  The  solutions  of  salts  are  of  different 
strengths,  but  the  great  majority  are  of  normal  (N),  or  even  smaller 
concentrations.  In  practice  they,  also,  are  still  further  consider- 
ably diluted  before  use.  If,  therefore,  we  give  the  fractions  ionized 
(total  molecules  of  ionogen  =  1)  in  decinormal  solutions  (except 
where  otherwise  specified),  the  reader  will  be  able  to  estimate  roughly 
the  proportion  of  each  kind  of  ions  in  any  application  of  the  reagent, 


IONIZATION 


367 


In  the  case  of  acids  containing  more  than  one  displaceable  hydrogen 
unit,  the  kind  of  ionization  on  which  the  figure  is  based  is  indicated  by 
a  period.  Thus  H.HCOa  means  that  the  whole  of  the  ionization  is 
assumed  to  be  into  H+  and  HCOs". 

FRACTION  IONIZED  IN  0.1  AT'  SOLUTIONS  AT  18° 
ACIDS 


Nitric  acid    ........  0.92 

Nitric  acid  (cone.,  62%)     .    .  0.09 

Hydrochloric  acid    .....  0.92 

Hydrochloric  acid  (cone.,  35%)  0  .  13 

Sulphuric  acid,  H.H.SO4     .    .  0.61 

Sulphuric  acid  (cone.,  95%)  .  0.01 

Hydrofluoric  acid     .....  0.15 

Oxalic  acid,  H.HC2O4      ...  0.50 

Tartaric  acid,  H.HT    ....  0.08 

Acetic  acid  (AT)    ......  0.004 

Acetic  acid    ........  0.013 


Carbonic  acid  (AT/25)     .    .    .  0.0021 

Hydrogen  sulphide,  H.HS      .  0.0007 

Boric  acid,  H.H2BO3       .    .    .  0.0001 

Hydrocyanic  acid    .....  0  .  0001 

Permanganic  acid  (AT/2)     .    .  0.93 

Hydriodic  acid  (AT/2)      ...  0.90 

Hydrobromic  acid  (A^/2)    .    .  0.90 

Perchloric  acid  (AT/2)      ...  0.88 

Chloric  acid  (AT/2)      ....  0.88 

Hypochlorous  acid      ....  0.0002 

Phosphoric  acid,  H.H2PO4     .  0.27 


Carbonic  acid,  H.HCO3 


0.0017      Water    ..........     0.061 


BASES 


Potassium  hydroxide      ...  0.91 

Sodium  hydroxide    .....  0.91 

Barium  hydroxide    .....  0  .  77 

Lithium  hydroxide  (AT)       .    .  0.63 
Tetramethylammonium     hy- 

droxide  (AT/16)     .....  0.96 


Ammonium  hydroxide    .    .    .  0.013 

Strontium  hydroxide  (AT/64)  0.93 

Barium  hydroxide  (AT/64)      .  0  .  92 

Calcium  hydroxide  (AT/64)     .  0.90 

Silver  hydroxide  (N/  1783)     .  0.39 

Water    .........   .  0.0«1 


SALTS 


Potassium  chloride      ....  0.86 

Potassium  nitrate    .....  0.83 

Potassium  acetate   .....  0.83 

Potassium  sulphate     ....  0.72 

Potassium  carbonate  ....  (0.71) 

Potassium  chlorate      ....  0.83 

Ammonium  chloride    ....  0.85 

Sodium  chloride  (A7)    ....  0.66 

Sodium  chloride  (N/2)    ...  0.74 

Sodium  chloride  ......  0.84 

Sodium  nitrate     ......  0.83 

Sodium  acetate    ......  0.79 

Sodium  sulphate      .   .   ,   ,   ,  Q.7Q 


Sodium  bicarbonate, 

Na.HCO3      .......  0.78 

Sodium  phosphate,  Na2.HPO4  0.73 

Sodium  tartrate   ......  0.69 

Barium  chloride       .....  0.77 

Calcium  sulphate  (AT/  100)     .  0.64 

Cupric  sulphate    ......  0.39 

Silver  nitrate    .......  0.81 

Zinc  sulphate    .....    .    .  0.40 

Zinc  chloride     .......  0.73 

Mercuric  chloride        .    .    .    (<0.01) 

Mercuric  cyanide     ,    ,   .   ,   .  Minute, 


368  INORGANIC  CHEMISTRY 

In  addition  to  their  use  in  showing  the  nature  of  the  reagents 
employed  in  the  laboratory  (p.  365),  these  numbers  show  also  to 
what  extent  any  pair  of  ionic  substances  will  unite  when  mixed  (see 
pp.  377,  380-382),  and  they  likewise  indicate  the  chemical  activity  of 
the  ionogens  when  in  solution  (see  next  section). 

Misapprehension  easily  arises  in  regard  to  the  inferences  that  may  be  drawn 
from  a  conductivity  value.  A  single  such  value,  say  that  for  salt  at  10  1.  dilution 
(92.5),  gives  no  information  about  the  extent  of  ionization.  We  must  have  the 
value  at  infinite  dilution  as  well,  that  is,  we  must  have  the  other  term  of  the  ratio 
corresponding  to  complete  ionization,  before  the  proportion  of  the  molecules 
ionized  at  the  10  1.  dilution  can  be  known.  Further,  we  must  have  the  values  of 
both  for  the  same  salt,  at  the  same  temperature  and  in  the  same  solvent,  for  the 
values  at  all  dilutions  change  markedly  when  any  one  of  these  conditions  is 
altered.  Thus  the  conductivity  of  normal  sodium  chloride  solution  at  50°  is  120, 
and  is  therefore  actually  greater  than  at  18°  when  the  dilution  is  infinite.  But 
at  50°  the  conductivity  at  infinite  dilution  is  185,  so  that  at  this  temperature  the 
degree  of  ionization  is  iff  or  0.65,  about  the  same  as  at  18°.  On  the  other  hand, 
when  a  little  alcohol  is  added  to  the  aqueous  solution,  the  conductivities  all 
diminish.  But  that  at  infinite  dilution  diminishes  also,  so  that  the  proportion  of 
the  material  ionized  does  not  seem  to  be  greatly  affected.  The  chief  effect  cf 
raising  the  temperature  is  simply  to  diminish  the  friction  opposing  the  motion  of 
the  ions  and,  therefore,  to  increase  the  conductivity.  The  change  is  about  2  per 
cent  for  each  degree.  Addition  of  alcohol,  on  the  other  hand,  increases  the  friction 
and  diminishes  the  conductivity.  There  is,  however,  a  real,  though  usually 
smaller,  change  in  the  degree  of  ionization  with  change  in  temperature.  When 
the  temperature  is  raised,  the  fraction  ionized  increases  or  diminishes  according 
as  the  heat  of  ionization  is  negative  or  positive  (cf.  p.  305),  and  conversely  when 
the  temperature  is  lowered. 

Degree  of  lonisation  of  Water.  —  If  we  consider  a  liter  of 
water  as  a  normal  solution  in  which  18  g.  (one  mole)  represents  the 
solute  and  the  rest  stands  for  the  solvent,  the  conductivity  for  complete 
ionization  into  H+  and  OH~  would  be  318  +  174  =  492.  The  actual 
ionization  is  one  ten-millionth  part  of  this.  In  other  words,  there  is 
only  one  ten-millionth  of  1  g.  of  hydrogen-ion  and  the  same  fraction  of 
17  g.  of  hydroxide-ion  in  a  liter  of  water.  A  column  of  water  1  cm. 
long  (18°)  conducts  less  well  than  a  column  of  mercury  of  equal  cross- 
section  and  over  1,700,000  miles  in  length. 

With  some  substances,  as  the  temperature  is  continuously  raised, 
the  fraction  ionized  increases  to  a  maximum  and  then  diminishes. 
With  water,  the  maximum  is  reached  at  229°  (Arrhenius)  or  250-275° 
(A.  A,  Noyes),  Ammonium  hydroxide  and  acetic  acid  also  exhibit 


IONIZATION  369 

similar  maxima.  In  these  cases,  the  heat  of  ionization  changes  its 
sign,  being  at  first  negative,  and,  beyond  the  maximum,  positive  (cf. 
p.  305). 

Relation  of  Ionization  to  Chemical  Activity.  —  These  tables 
may  be  used  for  reference.  The  import  of  the  following  general 
statements,  drawn  from  the  tables,  should  be  memorized: 

1.  Salts,  with  the  exception  of  those  of  mercury,  are  all  well 
ionized.     In  actions  involving  their  ions,  salts  are  therefore  all  of 
the  same  order  of  activity,  for  a  dilute  solution  of  every  salt  contains 
a  large  amount  of  the  ionic  components. 

Occasionally,  the  differences  in  the  degrees  of  ionization  of  salts 
have  to  be  considered.  The  rule  is  simple.  Salts  with  two  univalent 
radicals  are  the  most  highly  ionized  (e.g.,  KC1  86  per  cent,  KNOs  83 
per  cent).  Those  with  one  bivalent  radical  are  less  ionized  (e.g., 
K2SO4  72  per  cent,  K2C03  71  per  cent).  Those  with  two  bivalent 
radicals  are  still  less  ionized  (e.g.,  CuSO4  39  per  cent,  ZnSO4  40  per 
cent).  Those  containing  one  trivalent  radical  are  less  ionized  than 
are  those  with  one  bivalent  radical. 

2.  Acids  show  the  most  extreme  differences  in  their  degrees  of 
ionization.     That  is  to  say  their  solutions  must  contain  very  differ- 
ent concentrations  of  hydrogen-ion.     Since  their  activity  as  acids 
depends  on  this  substance  (p.  340),  and  the  activity  of  a  substance 
is  proportional  to  its  concentration  (p.  292),  it  follows  that  acids 
will  show  very  great  differences  in  apparent  chemical  activity.     At 
this  point,  therefore,  we  emerge  from  semi-physical  discussion  of 
the  subject  and  reach  something  of  definite,  practical  application 
in  chemical  work. 

The  data  show  that  acids  may  be  divided  roughly  into  four 
classes  with  different  degrees  of  acidic  activity: 

(a)  The  ionization  in  decinormal  solution  exceeds  70  per  cent; 
e.g.,  nitric  acid  and  hydrochloric  acid.     These  are  the  acids  which 
are  chemically  most  active,  for  their  solutions  contain  relatively 
high  concentrations  of  hydrogen-ion. 

(b)  The  ionization  is  between  70  and  10  per  cent;  e.g.,  sulphuric 
acid  and  phosphoric  acid.     These  acids  are  noticeably  less  active, 
for  their  solutions  contain  lower  concentrations  of  hydrogen-ion. 

(c)  The  ionization  is  between  10  and  1  per  cent;  e.g,,  acetic  acid. 
These  are  the  weaker  acids,  for  their  solutions  contain  very  small 
concentrations  of  hydrogen-ion. 

(d)  The  ionization  is  less  than  1  per  cent;   e.g.,  carbonic  and 


370  INORGANIC  CHEMISTRY 

boric  acids.  These  are  the  feeble  acids,  for  their  solutions  contain 
only  minute  concentrations  of  hydrogen-ion. 

3.  The  bases  show  two  classes: 

(a)  lonization  high;  e.g.,  potassium  hydroxide.  These  bases 
are  active,  for  their  solutions  contain  high  concentrations  of  hydrox- 
ide-ion. 

(6)  lonization  less  than  2  per  cent;  e.g.,  ammonium  hydroxide. 
These  bases  are  weak  on  account  of  the  low  concentration  of  hy- 
droxide-ion. 

4.  Water  is  less  ionized  than  any  other  substance  in  the  list.     It 
shows  therefore,  as  we  already  know,  usually  little  or  no  interaction 
with  acids,  bases,  or  salts,  and  hence  is  valuable  as  a  solvent  for 
these  substances.     Its  ions  are  H+  and  OH~,  and  it  is  thus  as  much 
an  acid  as  a  base. 

!  Exercises.  —  1.  With  solutions  of  the  following  substances, 
state,  (a)  what  will  be  the  products  of  electrolysis,  (b)  whether  each 
is  primary  or  secondary,  and  (c)  how  they  may  be  isolated  in  each 
case:  Potassium  chlorate,  potassium  iodide,  potassium  iodate,  silver 
sulphate,  sodium  peroxide,  sodium  fluoride. 

2.  Make  equations  (p.  358)  showing  the  ionic  and  molecular 
materials  in  solutions  of  potassium  bromide,  potassium  bromate, 
sodium  periodate,   aluminium  chloride,   zinc  sulphate.     Mark  the 
charges  on  the  ions  and  give  the  name  of  each  ionic  substance  (p. 
356). 

3.  Prepare  lists  of  other  anions  and  cations  which  have  been 
encountered,  giving  the  formula  and  number  of  charges  of  elec- 
tricity in  each  case. 

4.  How  many  coulombs  are  carried  by  or  will  deposit:  20  g.  of 
silver,  15  g.  of  antimony,  30  g.  of  chlorine,  60  g.  phosphate-ion  (PO4)? 

5.  What  current  strength  (in  amperes)  is  required  to  deposit: 
20  g.  of  silver  in  an  hour,  100  g.  of  iodine  in  5  minutes,  60  g.  of  anti- 
mony in  3  hours? 

6.  What  is  the  percentage  of  molecules  ionized  in:   deci-normal 
(N/ 10)  sodium  chloride,  centi-normal  (N / 100)  acetic  acid,  centi- 
normal  hydrochloric  acid  (p.  365)? 

7.  If  1  c.c.  of  dilute  hydrochloric  acid  (6JV)  is  added  to  30  c.c. 
of  an  aqueous  solution,  what  is  the  reacting  concentration  of  the  acid? 

8.  Classify  all  the  acids  in  the  table  (p.  367)  according  to  the 
four  classes  (p.  369). 

9.  Two  troughs,  one  4  inches  long  and  the  other  a  mile  lpng?  are 


IONIZATION  371 

filled  with  cupric  sulphate  solution,  and  a  plate  of  copper  with  a  wire 
connection  is  inserted  at  each  end  of  each  trough.  The  pairs  of 
plates  are  connected  simultaneously  with  a  battery  or  dynamo.  If 
there  is  any  difference,  in  which  cell  will  the  deposit  of  copper  on  the 
electrode  appear  first,  and  about  how  long  will  be  the  time  required 
for  its  appearance  in  the  other  cell? 


CHAPTER   XIX 
IONIC   SUBSTANCES   AND   THEIR  INTERACTIONS 

BEFORE  considering  the  typical  interactions  of  ionogens  in  solu- 
tion, we  must  have  a  clear  conception  of  the  peculiarities  of  these 
bodies  which  are  likely  to  affect  their  behavior.  The  facts  on  which 
such  a  conception  must  be  based  have  been  given  in  preceding 
chapters,  and  all  that  is  now  necessary  is  to  collect  and  apply  these 
facts. 

In  this  discussion,  after  enumerating  the  various  kinds  of  ionic 
substances,  it  must  be  made  clear  that  aqueous  solutions  of  ionogens 
are  mixtures  containing  several  solutes.  We  then  consider  the 
relations  of  the  ionic  and  the  molecular  substances,  in  equilibrium, 
when  a  single  ionogen  is  present.  It  must  also  be  shown  that  each 
kind  of  ions  is  a  distinct  substance  with  individual  physical  and 
chemical  properties.  Next,  salts  being  used  for  illustration,  the 
commonest  kind  of  interaction,  double  decomposition  between  iono- 
gens, will  be  discussed.  In  this  connection  precipitation  brings  up 
the  peculiar  state  of  equilibrium  between  the  undissolved  solute 
and  the  complex  of  molecules  and  ions  in  solution.  The  interaction 
of  acids  and  bases  (neutralization)  then  follows.  Weak  acids  and 
bases  and  hydrolysis  of  salts  are  next  discussed.  Finally,  the  five 
varieties  of  ionic  chemical  change  are  given,  and  the  practical  impor- 
tance of  actions  in  which  ions  play  no  part  is  emphasized. 

The  discussion  of  systems  in  equilibrium  in  the  present  chapter 
will  be  almost  purely  qualitative.  The  quantitative  consideration 
of  ionic  equilibria  (cf.  p.  359)  is  postponed  until  the  study  of  the 
metals  and  their  compounds  is  taken  up  (see  Chap.  XXXIV). 

The  Classes  of  Ionogens.  —  Acids  are  classified  according  to 
the  number  of  replaceable  hydrogen  units  in  their  molecules. 
Thus  chloric  acid  HClOs  is  a  monobasic  acid,  sulphuric  acid 
H2S04,  a  dibasic  acid  and  phosphoric  acid  H3PO4  a  tribasic  acid. 
These  terms  relate  to  the  fact  that,  in  neutralization  (see  p.  386) 
the  acids  interact  with  one,  two,  or  three  molecules  of  a  base  like 
sodium  hydroxide. 

372 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          373 

Bases  are  named  in  a  similar  way:  sodium  hydroxide  NaOH  is  a 
monoacid  base,  calcium  hydroxide  Ca(OH)2  is  a  diacid  base. 

Salts  like  KC1  and  Na2CO3  are  neutral  (see  acid  salts,  below)  or 
normal  salts,  and  NaKC03  and  Ca(OCl)Cl  (bleaching  powder)  are 
mixed  salts. 

The  most  interesting  classes  of  mixed  salts  are  the  acid  salts 
(p.  215)  and  the  basic  salts.  In  acid  salts,  like  NaHSO4  (p.  206)  and 
NaH2PO4  (p.  207),  all  the  hydrogen  of  the  acid  has  not  been  replaced 
by  a  metal.  In  basic  salts,  like  Ca(OH)Cl,  part  of  the  basic  hydroxyl 
remains. 

All  these  substances  are  ionogens  (p.  355).  The  mixed  salts  are, 
naturally,  dissociated  into  more  than  two  ionic  substances.  For  a 
fuller  discussion  of  these  and  other  salts  see  pp.  400-402. 

Ionic  Substances  Furnished  by  Acids.  —  The  mode  of  nam- 
ing ionic  substances  has  already  been  given  (p.  356). 

Acids,  e.g.,  HC1,  H2SO4,  when  dissolved  in  water,  all  furnish 
hydrogen-ion  H+  and  a  negative  ionic  substance  (anion),  e.g.,  Cl~, 
SO4=.  The  solutions  differ  from  those  of  salts  in  the  constant  pres- 
ence of  hydrogen-ion,  and  in  the  absence  of  any  other  positive  ion. 

Hydrogen-ion  H+  is  a  colorless  substance.  It  is  sour  in  taste, 
and  its  presence  is  recognized  by  the  fact  that  it  turns  blue  litmus 
red  (see  Indicators,  p.  391).  These  properties  serve  as  tests  for 
acids,  as  they  are  not  commonly  interfered  with  by  other  substances 
which  may  be  present.  Hydrogen-ion  is  univalent  and,  when 
combined  with  negative  radicals  of  salts,  gives  the  (molecular)  acids. 
The  activity  of  acids  depends  upon  the  concentration  of  the  hydrogen- 
ion  they  furnish  (p.  3C9),  and  therefore  upon  their  solubility  and 
the  degree  of  ionization  of  the  dissolved  molecules.  Some  furnish 
so  little  hydrogen-ion  that  their  action  on  litmus  can  hardly  be 
detected. 

The  substances  of  the  composition  HC1,  H2SO4,  and  so  forth, 
are  commonly  called  acids.  But  it  is  only  when  they  have  been 
dissolved  in  water,  or  some  other  ionizing  solvent,  that  they  show 
the  properties  characteristic  of  acids.  In  fact,  there  is  only  one  acid, 
hydrogen-ion  H+,  although  the  substances  which  give  it  by  dissocia- 
tion are  many.  The  parent  substances  are  salts  of  hydrogen,  in 
which  the  element  hydrogen  plays  the  part  of  a  metallic  element. 

Ionic  Substances  Furnished  by  Bases.  —  Bases,  e.g.,  KOH, 
NH4OH,  Zn(OH)2,  all  furnish  hydroxide-ion  OH~  and  some  positive 


374  INORGANIC  CHEMISTRY 

ionic  substance  (cation),  K+,  NH4+,  Zn"1"1".  Their  solutions  differ 
from  those  of  salts  in  the  constant  presence  of  hydroxide-ion  and 
in  the  absence  of  any  other  anion.  The  more  active  bases,  that  is, 
those  which  are  soluble  and  highly  dissociated,  so  that  they  give  a 
high  concentration  of  hydroxide-ion,  are  called  alkalies.  Such  are 
potassium  and  sodium  hydroxides.  They  are  often  named  caustic 
alkalies  and,  individually,  caustic  potash  and  caustic  soda.  The 
solutions  are  called  lyes. 

Hydroxide-ion  OH~  is  a  colorless  substance.  Properties  which 
serve  as  tests  for  bases  are  that  hydroxide-ion  possesses  a  soapy 
taste  and  turns  red  litmus  blue  (see  Indicators,  p.  391).  It  is  uni- 
valent,  and  combines  with  positive  radicals  to  form  (molecular)  bases. 

The  common  bases,  with  the  exception  of  the  hydroxides  of 
potassium,  sodium,  barium,  strontium,  calcium,  and  ammonium,  are 
but  slightly  soluble  in  water.  Hence,  zinc  hydroxide,  for  example, 
although  it  dissolves  sufficiently  to  enable  chemical  action  to  take 
place  slowly,  does  not  give  enough  hydroxide-ion  at  one  time  to  affect 
litmus  paper.  Magnesium  hydroxide  and  lead  hydroxide  turn  red 
litmus  paper  blue  with  difficulty.  Doubtless  the  few  molecules  that 
do  dissolve  are  almost  all  ionized : 

Mg(OH)2  (solid)  ±?  Mg(OH)2  (dslvd)  ±5  Mg+++  20IT 

but  all  the  dissolved  materials  put  together  (0.01  g.  per  1.)  will  scarcely 
be  weighable  unless  a  considerable  volume  of  the  solution  is  evap- 
orated. 

Substances  like  potassium  hydroxide,  ammonium  hydroxide 
NH4OH,  and  zinc  hydroxide  Zn(OH)2,  are  commonly  called  bases. 
But  it  is  only  in  their  aqueous  solutions  that  the  basic  properties 
appear.  There  is  only  one  base,  namely,  hydroxide-ion  OH~,  and 
these  substances  are  simply  the  source  of  it.  The  parent  substances 
are  salts  of  some  metallic  element,  or  group  playing  the  part  of  a 
metallic  element  (e.g.,  NHi),  in  which  hydroxyl  is  the  negative  radical. 

The  name  "base"  was  originally  applied  to  the  non-volatile,  and  therefore 
seemingly  more  fundamental  part  of  a  salt  that  remained  behind  when  the  salt 
was  heated.  Usually  the  negative  radical  is  disintegrated,  as  in  heating  calcium 
carbonate:  CaCO3  — »  CaO  +  CO2 1  .  But,  as  a  matter  of  fact,  it  is  generally 
the  oxide  and  not  the  hydroxide  of  the  metal  that  remains.  Still,  the  oxide,  for- 
merly named  tne  base,  often  readily  gives  the  hydroxide  (cf.  p.  149)  to  which  the 
term  "base"  is  now  applied,  and  behaves  similarly  to  it  in  many  interactions 
(cf.  p.  213). 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          375 

Ionic  Substances  Furnished  by  Salts.  —  Salts  furnish  posi- 
tive and  negative  ionic  substances,  which  may  be  either  simple 
or  composite,  Na.Cl,  Na.NO3,  NH4.C1,  NH4.NO3.  Some  ionic  sub- 
stances are  colored,  Cu**  (cupric-ion)  blue,  Cr"1"1"1"  reddish-violet, 
Co++  pink,  MnC>4~  (permanganate-ion)  purple,  Cr2C>7=  (dichromate- 
ion)  orange,  but  most  of  them  are  colorless,  K+,  Na+,  Zn"1"1",  Cl~~, 
I~,  NO3~~,  SC>4=.  They  vary  in  taste,  some  being  salt,  some  astringent, 
some  bitter.  They  carry  electricity,  but  relatively  less  well  than 
do  hydrogen-ion  and  hydroxide-ion,  on  account  of  their  slower 
migration.  The  ionic  materials  characteristic  of  salts  do  not  effect 
litmus,  and  individual  tests  are  required  for  each.  Usually  we  add 
some  other  ionic  substance,  with  which  the  ion  thought  to  be  present 
combines  to  form  an  insoluble,  molecular  substance  of  known  color, 
or  appearance,  and  examine  the  precipitate  if  any  appears.  Thus, 
when  the  presence  of  chloride-ion  Cl~  is  suspected,  we  may  add  a 
solution  containing  silver-ion  Ag+,  expecting  to  obtain  a  precipitate 
of  silver  chloride  AgCl  (Cl~  +  Ag+  — >  AgCl  JJ.  -In  dilute  solutions 
of  salts,  the  ions  are  almost  always  numerous  in  comparison  with 
the  molecules  (p.  369),  so  that  salts  are  practically  all  active  and 
their  solutions  almost  always  respond  readily  to  the  tests  for  the  ions 
they  contain.  The  art  of  detecting  the  various  ionic  substances 
present  in  a  solution  constitutes  a  large  part  of  the  branch  of  chem- 
istry called  qualitative  analysis  (see  below). 

All  the  known  ionic  substances  are  found  in  solutions  of  salts. 
The  only  ions  which  are  not  characteristic  of  salts,  although  some- 
times occurring  in  their  solutions  (see  acid  and  basic  salts,  above), 
are  hydrogen-ion  H+,  and  hydroxide-ion  OH~. 

It  will  assist  the  reader  if  the  following  facts  are  kept  in  mind. 
The  elements  which  can  form  a  simple  positive  ion  are  the  metallic 
elements  (p.  214,  and  see  Chaps.  XXII  and  XXXII).  Non-metallic 
elements,  like  nitrogen,  may  be  present  in  a  positive  ion,  as  in  NEU*, 
but  never  exclusively.  In  other  words,  we  know  no  such  substances 
as  nitrogen  sulphate,  or  carbon  nitrate.  Conversely,  the  metals  are 
frequently  found  in  the  negative  ion,  but  never  constitute  it  exclu- 
sively. They  are  then  usually  associated  with  oxygen,  as  in  MnO-T, 
and  Cr2O7=. 

Solutions  of  lonogens  are  Mixtures,  in  which  Each  Kind 
of  Ion  Acts  Independently.  —  We  are  accustomed  to  regard  a 
bottle  of  sodium  chloride  solution  as  containing  but  one  thing,  aside 
from  the  water.  We  must  now  think  of  it  as  containing  at  least  three 


376  INORGANIC   CHEMISTRY 

dissolved  substances,  any  one  of  which  might  be  alone  responsible 
for  some  property  of  the  solution.  The  same  idea  must  accompany 
our  use  of  every  solution  of  an  ionogen. 

Numberless  facts  show  that  each  kind  of  ion,  for  example  cupric-ion, 
has  an  individual  set  of  physical  and  chemical  properties  and  behaves 
in  many  ways  as  if  alone  present  in  the  solution.  We  shall  meet  with 
much  evidence  of  this  in  the  sequel.  Some  facts  tending  to  prove  it, 
that  have  already  been  given,  may  be  recalled  (cf.  p.  339). 

If,  in  comparing  the  migration  speeds  of  any  element,  say  copper, 
in  different  salts  (p.  345),  they  were  the  motions  of  substances  like 
Cu(NOs)2,  CuBr2,  CuS04,  that  we  were  comparing,  all  analogy  teaches 
us  that  the  speeds  with  which  they  would  move  should  vary  widely. 
That  the  blue  color  drifts  always  at  the  same  pace  shows  that  it  is  the 
same  substance,  namely  cupric-ion  Cu++,  that  we  are  observing. 

If,  in  solutions  of  the  different  permanganates,  KMnO4,  NaMnO4, 
Ba(MnO4)2,  and  so  forth,  the  dissolved  bodies  were  different  in  each 
case,  we  should  confidently  expect  the  purple  colors  of  the  solutions  to 
differ  markedly  in  shade.  But,  for  dilute  solutions  of  equivalent 
concentrations,  when  strict  examination  is  made,  the  tints  are  found 
to  be  absolutely  identical.  We  are  therefore  simply  comparing 
different  mixtures  all  containing  the  same  proportion  of  the  same 
free,  colored  body,  MnO4~. 

In  phosphorus  pentachloride  vapor  (p.  260),  the  fully  liberated 
trichloride  and  chlorine  are  prominent  components.  Diminishing 
the  volume  of  a  fixed  amount  of  this  mixture,  by  compression,  throws 
more  chlorine  into  combination  and  the  total  absorption  of  blue 
light  (from  which  the  greenish-yellow  color  is  derived)  becomes  less, 
the  compounds  of  phosphorus  being  both  colorless.  Increasing  the 
volume,  on  the  other  hand,  promotes  the  dissociation  and  increases 
',he  depth  of  the  yellow  color.  The  system  of  ions  and  molecules 
in  equilibrium  in  a  solution  of  cupric  bromide  (see  below),  or  any 
other  ionogen,  behaves  in  exactly  the  same  way.  The  components 
possess  and  exhibit  individual  properties,  much  like  the  components 
of  a  gaseous  mixture  (p  111),  both  in  this  and  in  other  respects. 

All  solutions  of  acids  are  sour  in  taste,  irrespective  of  the  nature 
of  the  negative  ion,  while  salts  containing  the  same  negative  radical 
are  not  sour  at  all.  Hence  in  solutions  of  acids  we  are  tasting  the 
same  free  substance,  hydrogen-ion  H+.  Similarly,  in  solutions  of  all 
alkalies,  we  note  the  soapy  taste  of  hydroxide-ion  OH~. 

These  illustrations  concern  physical  properties.  In  the  later 
sections  we  shall  learn  that  an  ionic  material,  such  as  bromide-ion 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          377 

or  cupric-ion,  has  specific  chemical  properties  irrespective  of  the 
nature  of  its  concomitants. 

IONIC    EQUILIBRIUM 

The  Ionic  Equilibrium  Involving  a  Single  lonogen.  —  In 

the  ionization  of  a  molecular  substance,  the  chemical  change  is 
incomplete  and  the  system  reaches  a  condition  of  equilibrium  (p.  358). 
The  action  is,  therefore,  reversible,  and  there  are  thus  two  routes  to 
the  same  equilibrium  point.  This  fact  must  not  be  forgotten,  for 
we  have  to  consider  the  union  of  ionic  substances  even  more  often 
than  the  converse  change.  Now,  the  degrees  of  ionization  of  various 
ionogens  tell  us  the  location  of  the  equilibrium  point,  and  therefore 
the  extent  of  the  chemical  change  involved  in  reaching  this  point  by 
either  route,  that  is,  either  by  the  dissociation  of  molecules  or  by  the 
union  of  ions.  In  a  class  of  interactions,  of  which  all  are  incom- 
plete, and  only  those  are  interesting  and  useful  which  approach 
completeness,  we  require  some  means  of  knowing  which  are  com- 
plete and  why  they  are  so.  The  table  of  fractions  ionized  (p.  367) 
supplies  most  of  the  required  information. 

To  illustrate,  take  the  case  of  a  single  ionogen.  When  we  place 
hydrogen  chloride  in  decinormal  solution,  0.92  of  the  molecules  dis- 
sociate. Conversely,  when  we  start  with  the  hydrogen-ion  and 
chloride-ion,  say  by  mixing  two  solutions,  each  of  which  contains 
one  of  them  (along  with  another  ion), -then  1  —  0.92,  or  only  0.08 
(=8  per  cent)  of  these  ionic  substances  will  combine. 

This  exemplifies  the  case  of  an  active  acid.  The  following  equa- 
tions show  the  data  for  six  typical  substances  in  N/IQ  solution, 
namely,  two  acids,  two  bases,  and  two  salts: 

(8%)HC1    ^H+     +CF(92%),     (98.7%)  HC2H3O2^H+ +  C2H3Or(1.3%) 
(9%)KOH^K+     +OH~(91%),   (98.7%)  NH4OH   ^±NH4+  +  OH"  (1.3%) 
(16%)  NaCl  <=>  Na+  +  CP  (84%),          (61%)  CuSO4  *±  Cu^  +  SO4=  (39%) 

These  samples  are  chosen  to  illustrate,  in  each  pair,  the  extremes. 
Thus,  when  potassium-ion  and  hydroxide-ion  are  brought  together 
little  union  takes  place,  while  with  ammonium-ion  and  hydroxide- 
ion  the  union  is  practically  complete.  In  the  case  of  the  soluble 
salts,  however,  there  are  almost  (p.  369)  no  cases  of  considerable 
union  of  the  ions  in  dilute  solutions.  The  case  of  water,  on  the 
other  hand,  is  one  of  the  most  extreme: 

(99.96%)  H20  <=±  H+  +  OH-  (0.041%). 
Hydroxide-ion  and  hydrogen-ion  thus  unite  almost  completely. 


378  INORGANIC  CHEMISTRY 

Similar  reasoning  enables  us  to  handle  the  more  complex,  but 
very  common  case  of  the  mixing  of  two  ionogens.  The  degrees  of 
ionization  tell  us  the  exact  condition  of  each  system  separately, 
before  mixing.  The  result  of  the  mixing  is  best  understood  by 
viewing  the  change  as  consisting  in  a  displacement  of  each  of  the 
equilibria  by  the  action  of  the  components  of  the  other.  We  con- 
sider, therefore,  next,  the  displacement  of  ionic  equilibria. 

The  Displacement  of  Ionic  Equilibria.  —  Equilibria  are  dis- 
placed by  changes  which  favor  or  disfavor  one  of  the  opposed  actions 
(p.  290).  There  may  be  either,  (1)  a  physical  change  in  the  condi- 
tions, or  a  chemical  interaction  which  (2)  increases  the  amount  of, 
or  (3)  removes  one  of  the  interacting  substances.  Each  of  these 
may  be  illustrated  in  turn. 

1.  As  d,n  example  of  the  first,  we  have  the  effect  of  changing  the 
amount  of  the  solvent  (p.  337).  Adding  more  of  the  solvent  reduces 
the  concentration  of  the  ionic  materials  and  disfavors  their  union, 
so  that  it  indirectly  promotes  dissociation.  The  larger  the  volume 
in  which  the  ions  are  scattered,  the  less  often  will  they  meet,  and 
the  smaller  the  amount  of  combination.  On  the  other  hand, 
evaporating  off  a  part  of  the  solvent  favors  the  encounters  of  the 
ions  and  promotes  combination.  When  the  solvent  is  at  last  entirely 
gone,  the  whole  material  is  molecular. 

In  cases  where  the  ionic  and  molecular  substances  are  all  color- 
less, these  changes  can  be  followed  only  by  a  study  of  the  freezing- 
points  or  other  similar  properties  of  the  solutions  (p.  336).  But 
when  the  substances  are  of  different  colors,  f  he  changes  can  also  be 
seen.  Thus,  cupric  bromide  in  the  solid  form  is  a  jet  black,  shining, 
crystalline  substance.  When  treated  with  a  small  amount  of  water 
it  forms  a  solution  which  is  of  a  deep  reddish-brown  tint,  giving  no 
hint  of  resemblance  to  a  solution  of  any  cupric  salt.  This  doubtless 
represents  the  color  of  the  molecules.  When  more  water  is  added,  the 
deep  brown  gives  place  gradually  to  green,  and  finally  to  blue.  The 
latter  is  the  color  of  the  cupric-ion  Cu"1"1",  and  is  familiar  in  all  solutions 
of  cupric  salts.  The  colorless  nature  of  solutions  of  potassium  and 
sodium  bromides  shows  that  bromide-ion  Br~  is  without  color. 
Hence,  in  the  present  instance  it  is  invisible.  We  are  thus  watching 
the  forward  displacement  of  the  equilibrium: 

CuBr2  (brown)  *=t  Cu++  (blue)  +  2Br~. 
If  1  g.  of  the  solid  is  taken,  it  dissolves  in  about  its  own  weight  of 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          379 

water,  and  independent  measurement  shows  that  there  is  relatively 
little  ionization.  Hence  the  solution  is  deep  brown.  When  10  c.c. 
of  water  has  been  added,  70  per  cent  of  the  salt  is  ionized,  and  the 
solution  is  green.  With  40  c.c.  of  water,  only  19  per  cent  remains 
in  molecular  form,  and  the  blue  color  of  the  cupric-ion  entirely 
overbears  the  tint  of  the  molecules.  If  we  now  remove  the  water 
by  evaporation,  all  these  changes  are  reversed.  When  30  c.c.  of 
the  water  has  been  driven  off,  the  solution  is  green.  As  the  evap- 
oration of  the  remaining  10  c.c  progresses,  the  brown  color  appears. 
When  the  water  is  all  gone,  the  black  residue  remains.  Here  we 
are  observing  the  backward  displacement  of  the  equilibrium, 
CuBr2  ±=>  CU++  +  2Br~. 

2.  Cupric  bromide  may  be  used  to  illustrate  also  the  chemical 
methods  of  displacing  equilibria.     Thus,  we  may  show  the  effect  of 
adding  more  of  one  of  the  reacting  substances.     If,  at  the  green  stage, 
we  dissolve  solid  potassium  bromide  in  the  liquid  (KBr  *=±  K+  +  Br~), 
the  increased  concentration  of  bromide-ion  causes  more  extensive 
interaction  of  the  ions,  and  the  molecules,  with  their  brown  color, 
become   prominent   again.     Adding   cupric    chloride   increases   the 
concentration  of   cupric-ion  and  has  the  same  effect.     In  either 
case,  renewed  dilution  with  water  reduces  the  concentrations  of  all 
the  ions  once  more,  the  molecules  become  fewer,  and  the  brown 
color  is  displaced  by  the  blue  for  the  second  time. 

3.  Finally,  the  displacement  of  the  same  equilibrium  by  remov- 
ing one  of  the  interacting  substances  may  be  illustrated.     Thus,  if 
the  chocolate-brown  solution,  in  which  molecular  cupric  bromide 
predominates,  is  shaken  with  pulverized  lead  nitrate  (and  filtered), 
two  changes  are  noticed.     A  pale  yellow  precipitate  of  lead  bromide 
appears  (Pb"1"1"  +  2Br~  — » PbBr2  j  ),  and  the  brown  color  fades  into 
green.     Here  the  displacement  is  the  .opposite  of  the  last.     Instead 
of  reinforcing  one  of  the  ions,  we  have  reduced  the  concentration, 
and  in  fact  almost  entirely  removed  one  of  them,   namely  Br~~. 
This  has,  naturally,  stopped  the  interaction  of  the  Cu++  and  Br~ 
which  reproduces  the  brown,  molecular  CuBr2.     Hence  the  disso- 
ciation of  the  latter  has    continued  to  exhaustion    of  the  whole 
molecular  material. 

The  reader  will  find  that  the  behavior  of  these  ionic  equilibria, 
and  the  way  in  which  we  discuss  and  explain  it,  are  complete  parallels 
of  the  behavior  and  explanation  in  the  case  of  ordinary  equilibria 
(pp.  169,  301),  which  should  now  be  reexamined.  The  illustrations 
in  the  present  section,  and  particularly  the  third  (cf.  p.  208),  should 


380  INORGANIC  CHEMISTRY 

be  considered  until  every  feature  is  perfectly  clear.  They  furnish 
the  key  to  understanding  the  applications  which  follow.  One  fact 
must  not  escape  notice,  and  that  is  that  in  none  of  the  three  instances 
was  the  forward  action  (the  dissociation)  in  itself  affected.  The 

molecules  of  cupric  bromide  have,  as  we  should  expect,  a  certain 
tendency  to  decompose.  No  encounters  between  these  molecules 
are  required  for  mere  decomposition.  Hence  their  decomposition  is 
not  influenced  by  their  nearness  to,  or  remoteness  from  one  another 
(illustration  1),  nor  by  the  presence  of  any  other  kinds  of  molecules 
or  ions  (illustrations  2  and  3).  The  effect,  whether  it  involved  an 
apparent  increase,  or  a  diminution  of  the  dissociation,  was  always 
accomplished  by  altering  the  concentration  of  the  ionic  substances,  and 
therefore  the  extent  of  the  reverse  action. 

DOUBLE  DECOMPOSITION  OF  Two  SALTS.     PRECIPITATION. 
CHEMICAL  PROPERTIES  OF  IONIC  SUBSTANCES 

Applications:     Double   Decomposition   in   Solution.  —  We 

are  now  prepared  to  consider  the  general  case  of  mixing  the  solu- 
tions of  two  ionogens. 

When  solutions  of  two  ionized  substances  are  mixed,  the  first 
reflection  which  occurs  to  us  is  that  each  of  these  has  been  diluted 
by  the  water  in  which  the  other  was  dissolved,  so  that  the  first 
effect  will  be  to  increase  the  degree  of  ionization  of  both  to  a  certain 
extent. 

The  next  consideration  is,  however,  that  we  have  produced  a 
mixture  of  four  ions,  which  must  have  at  least  some  tendency  to 
unite  crosswise.  Thus  potassium  chloride  and  sodium  nitrate  in 
dilute  solution  are  very  greatly  ionized  before  mixing.  The  re- 
versible actions,  represented  by  the  horizontal  pair  of  the  following 
equations,  have  taken  place  extensively.  But,  by  mixing  the 
liquids,  we  have  brought  into  presence  of  -^.p,  <_  -^-+  ,  ~,_ 

one  another  two  new  pairs  of  positive  and     AT  ATr.   "  ~^_  , Trk  _  .   ,T   , 
TT  JN  aJN  Oz  £=*  JN  Us  +  JN  a+ 

negative  ions.     Hence,  two  other  reversi-  i 

ble  actions,  the  vertical  ones,  will  be  set 

.,      .,',  -i         .-.       /.      i  «i»i 

up  and  will  proceed  until  a  fresh  equilib- 
rium of  all  the  ions  with  all  four  kinds  of  molecules  has  been  reached. 
Thus  far  the  description  will  fit  any  case  of  mixing  solutions  of  two 
ionogens. 

Now,  in  this  particular  instance,  what  is  the  actual  extent  of  such 
interaction  as  has  occurred?  To  answer  this  question  we  require  to 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          381 

know  the  proportion  of  molecules  to  ions  in  a  solution  of  each  of  the 
four  salts  (p.  367).  In  decinormal  solutions  it  is  KC1,  14:86; 
NaN03,  17:83;  KN03,  17:83,  NaCl,  16:84,  so  that  the  salts 
are  all  equally  well  ionized.  It  is  a  good  plan  to  add  these  pro- 
portions in  the  formulation,  and  to  insert  heavy  arrows  to  indicate 
the  preponderating  direction  in  each  equilibrium.  Furthermore, 
in  a  diluted  mixture,  such  as  we  are  considering,  the  proportions  of 
ions  are  greater  than  these  figures  indicate.  Hence,  practically  no 
chemical  action  has  occurred. 


(14%)KClfc?K-1-  +     01-  (86%) 

(17%)  NaNO3  *=>  NOr  +     Na+  (83%) 

It  IT 

KNO3  NaCl 


That  this  inference  is  correct  is  shown  by  independent  evidence. 
Thus  when  the  solutions  of  salts  are  mixed,  and  no  precipitation  oc- 
curs, no  thermal  effect  is  observable.  This  fact  has  been  known  since 
1842  as  Hess'  law  of  thermoneutrality.  There  is  also  no  change  in 
color  or  in  volume  on  mixing  such  solutions.  Again,  if  the  solutions 
(about  N/4i)  are  placed  in  a  cell  (Fig.  106,  p.  363),  so  that  the  one  forms 
a  layer  below  the  other  (the  solution  to  be  used  for  the  lower  layer 
is  weighted  with  sugar),  no  change  in  conductivity  is  noticed  when 
the  solutions  are  stirred  together.  Hence  no  change  in  the  number 
of  ions  has  occurred. 

We  conclude,  then,  that  when  two  highly  ionized  substances  are 
mixed,  and  the  possible  products  are  also  highly  ionized,  soluble 
substances,  then  practically  no  chemical  action  occurs.  This  rule  ap- 
plies to  all  soluble  salts  (p.  369)  and  to  mixing  salts  with  the  highly 
ionized  acids  or  bases. 

In  view  of  the  above  explanation,  the  old  question  of  whether  such  a  solution 
contains  the  first  pair  of  salts,  or  the  second  pair,  represented  in  the  double  de- 
composition, KC1  +  NaNO3  «=*  KNO3  +  NaCl,  loses  its  whole  point.  The  solu- 
tion contains  neither  the  initial  molecular  substances  nor  the  molecular  products, 
in  appreciable  amount. 

Conversely,  when  two  ionized  substances  are  mixed,  an  extensive 
chemical  change  does  ensue  in  two  cases,  namely: 

1.  When  one  of  the  possible  products  is  an  insoluble  substance 
and  precipitation  occurs,  for  this  removes  the  ions  used  in  forming 
the  insoluble  body. 


382  INORGANIC   CHEMISTRY 

2.  When  one  of  the  possible  products,  although  soluble,  is  little 
ionized,  as  in  neutralization,  for  this  likewise  removes  the  ions  re- 
quired to  form  molecules  of  the  product.  We  proceed,  therefore, 
to  discuss  these  two  important  classes  of  actions. 

Precipitation.  —  A  typical  case  of  precipitation  occurs  when 
we  mix  dilute  solutions  of  silver  nitrate  and  sodium  chloride. 


(16%)  NaCl  fc?  Na+    +  Cl~  (84%) 
(19%)  AgN03  ^  N03~  +  Ag+  (81%) 

IT         IT 

NaNO3    AgCl(dslvd) 
«17%)       IT 

AgCl  (solid) 

Here,  since  the  four  substances  are  all  salts,  they  are  all  highly 
ionized.  If  they  were  all  soluble,  then,  in  dilute  solutions,  perhaps 
5  per  cent  of  each  salt  would  be  in  molecules  and  the  rest  in  ionic 
form.  But  the  molecules  of  silver  chloride  are  excessively  insoluble. 
In  all  cases  of  precipitation,  we  look  up  the  solubilities  of  the  possible 
products  (see  Table  of  Solubilities,  inside  the  front  cover).  Here  we 
find  that  one  liter  of  water  will  dissolve  only  0.0016  g.  silver  chloride 
(this  quantity  includes  both  ions  and  molecules).  So  the  concentra- 
tion of  the  AgCl  (dslvd)  becomes  almost  zero  through  precipitation. 
So  far  as  it  is  in  solution,  however,  being  a  salt  and  very  dilute,  it 
is  practically  all  ionized.  The  precipitation  displaces  the  equilibrium, 
for,  the  dissociation  having  thus  ceased,  those  of  the  ions  Ag+  and 
Cl~  which  combine  are  not  replaced  by  others.  Hence  the  silver- 
ion  and  chloride-ion  almost  disappear.  This  occurrence  affects  in 
turn  the  equilibria  with  Na+  and  NO3~,  so  that  the  NaCl  and  AgNO3 
become  completely  ionized.  Hence  the  concentrations  of  NaCl 
and  AgNO3,  of  Ag+  and  Cl~,  and  of  the  dissolved  AgCl,  all  become 
practically  zero  at  last.  The  system  finally  contains  only  a  precipi- 
tate of  molecular,  solid  silver  chloride  and  a  solution  of  the  three 
substances,  Na+  +  NO3~  ±5  NaNO3,  in  equilibrium.  By  far  the 
greater  part  of  this  material  in  solution  is  the  ionic,  namely  the  Na+ 
and  the  NO8~. 

It  will  be  noted  that  precipitation  concerns  only  the  molecules, 
directly,  and  that  the  ions,  if  there  are  any,  are  involved  only  in- 
directly. The  ions  are  in  equilibrium  with  the  dissolved  molecules, 
not  with  the  precipitate. 

AB  (solid)  <±  AB  (dslvd)  i±  A+  +  B~ 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          383 

Hence,  when  a  substance  dissolves,  it  does  so  in  molecular  form,  and 
ions  are  subsequently  generated  from  some  of  these  molecules  until 
equilibrium  is  reached.  Conversely,  when  molecules  come  out  of 
solution,  as  the  result  of  cooling  or  precipitation,  the  diminished 
concentration  of  the  dissolved  molecules  enables  more  ions  progres- 
sively to  unite  until  the  whole  system  has  adjusted  itself  to  the  new 
conditions. 

To  avoid  a  misconception,  note  that  the  answer  to  the  question, 
"Is  silver  chloride  a  highly  ionized  substance?"  is  "Yes."  Since 
it  is  a  salt,  we  expect  this.  True,  very  little  of  it  dissolves,  so  that 
it  cannot  give  many  ions  to  a  solution.  But  little  or  much  ionized 
refers  to  the  proportion  ionized  of  the  material  which  is  in  solution. 
With  undissolved  material  ionization  has  nothing  to  do. 

It  should  be  noted  that,  when  the  solutions  are  mixed,  as  in  the 
foregoing  example,  strictly  speaking,  the  chief  interaction  taking 
place  is  the  production  of  the  insoluble  body.  The  largest  part  of 
the  chemical  action  may  be  formulated  thus: 

Ag+  +  Cl-  ->  AgCl. 

The  chief  change  that  has  as  yet  befallen  the  ions  of  sodium  nitrate 
is  that  they  have  been  transferred  from  two  separate  vessels  into 
one.  Potentially  the  salt  has  been  formed.  But  the  actual  union 
of  its  ions,  to  give  the  second  product  in  the  molecular  condition, 

Na+  +  N03-  -»  NaN03, 

comes  about  only  when,  at  some  subsequent  time,  if  at  all,  the 
water  is  evaporated  away. 

The  foregoing  formulation  and  explanation  apply  to  every  case 
of  mixing  ionogens  where  precipitation  occurs,  that  is,  where  the 
products  are  insoluble  acids,  bases  or  salts. 

If  the  least  soluble  of  the  four  salts  is  more  soluble  than  silver 
chloride,  more  concentrated  solutions  are  required  to  secure  precipi- 
tation. The  interaction  of  hydrogen  chloride  and  sodium-hydrogen 
sulphate  (p.  208)  is  of  this  nature: 

TTO1  *  —  TT"'"  -I-  f^l~~    } 
NaHS04  *  iSor  +  Na+  |  *  NaC1  (dslvd)  ^  NaC1 


Ionic  Double  Decomposition  and  Affinity.  —  It  is  quite 
clear  that  the  complete  formation  of  acids,  bases,  and  salts  by  pre- 
cipitation is  purely  a  result  of  mechanical  details  concerning  solu- 
bility, and  shows  nothing  about  the  degree  of  affinity  between  the 


384  INORGANIC  CHEMISTRY 

constituent  ions.  Again,  the  union  of  ions  to  form  feebly  ionized 
substances  only  shows  the  tendency  of  the  ionic  materials  to  unite, 
and  may  be  complete  where  the  free  elements  have  little  mutual 
affinity,  and  vice  versa.  Thus,  hydrogen-ion  and  hypochlorite-ion 
CIO  unite  almost  completely,  while  hydrogen-ion  and  chloride-ion 
hardly  unite  at  all.  Yet  hypochlorous  acid  HC1O  is  very  unstable, 
while  hydrogen  chloride  is  just  the  reverse.  Ionic  double  decomposi- 
tions, consequently,  give  no  clue  to  the  activities  of  the  free  materials. 

Individual,  Specific  Chemical  Properties  of  Each  Ionic 
Material.  —  We  wrote  the  equation  for  the  formation  of  silver 
chloride  Ag+  +  Cl~  — >  AgCl,  as  if  silver-ion  and  chloride-ion  were  the 
only  substances  concerned  in  the  action.  Further  study  shows  this 
to  be  justifiable.  Thus,  hydrochloric  acid,  cupric  chloride,  and 
dozens  of  other  chlorides  may  be  used  instead  of  sodium  chloride 
and  give  silver  chloride  just  as  readily.  The  sodium-ion  had  nothing 
to  do  with  the  result.  Of  course  we  cannot  get  a  solution  containing 
chloride-ion  alone.  Like  a  vessel  in  which  to  make  the  experiment, 
some  positive  ion  is  required.  But,  like  the  rest  of  the  apparatus, 
this  ion  may  be  varied  indefinitely,  is  not  altered  in  the  course  of  the 
change,  and  may  therefore  be  dispensed  with  in  the  equation.  The 
nitrate-ion  NO3~  which  accompanied  the  silver-ion  is  similarly  a 
part  of  the  apparatus,  for  silver  sulphate  solution  works  just  as 
well  as  silver  nitrate. 

That  chloride-ion  is  a  substance  with  specific  chemical  properties, 
is  easily  demonstrated.  It  forms  silver  chloride  whenever  it  en- 
counters silver-ion.  Other  substances,  even  when  they  contain 
chlorine,  lack  this  property.  Chloroform  CHCls  and  chlorobenzene 
CeH5Cl,  in  a  solvent  in  which  ionogens  are  dissociated,  do  not  interact 
when  silver  nitrate  is  added.  They  give  no  chloride-ion  and,  in  fact, 
remain  un-ionized.  Potassium  chlorate  KC103  and  perchlorate 
KCKV  and  chloracetic  acid  HCO2CH2C1,  with  silver-ion,  fail  like- 
wise to  give  silver  chloride.  They  are  ionized,  but  chloride-ion  is 
not  one  of  the  ions  of  any  of  them.  The  ions  ClOs~,  C1O4~,  and 
C02CH2C1",  have  properties  of  their  own,  and  their  compounds  with 
silver-ion  are  soluble. 

Other  chemical  properties  of  chloride-ion  are:  That  it  unites 
also  with  lead-ion  Pb++  and  mercurous-ion  Hg+,  forming  insoluble 
chlorides  (p.  226).  It  is  discharged  and  liberated  as  free  chlorine  by 
fluorine  (p.  281) : 

2H+  +  2Cr  +  F2  -»  2H+  +  C12  T 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          385 

Since  the  hydrogen-ion  is  not  affected  and  other  chlorides  behave  in  a 
similar  manner,  the  positive  ion  may  be  omitted :  2C1~  +  F£  — >  Cl2 
-f  2F~. 

Finally,  chloride-ion  has  relatively  little  tendency  to  unite  with 
other  ions,  or,  in  other  words,  the  compounds  of  chloride-ion  with 
most  other  ions  are  highly  ionized.  Thus  it  combines  with  hydrogen- 
ion  to  the  extent  of  only  8  per  cent  (p.  377)  in  decinormal  solution. 
In  this  respect  it  differs  markedly  from/ree  chlorine,  just  as  hydrogen- 
ion  differs  from  hydrogen.  The  free  elements  unite  with  vigor  and 
completely.  Hydrogen  chloride  is  easy  to  dissociate  into  ions, 
but  difficult  to  dissociate  into  its  constituent  elements.  Nothing 
could  show  more  strikingly  than  this  that  the  ionic  materials  have 
chemical  properties  of  their  own. 

Similarly,  barium  salts  and  ordinary  sulphates  give,  when  mixed, 
a  precipitate  of  barium  sulphate  BaSC>4.  Here  we  encounter  a 
property  of  barium-ion  Ba++  and  sulphate-ion  SC>4=.  But  potassium- 
ethyl  sulphate  KC2H5SO4,  in  spite  of  its  name,  will  not  give  this 
reaction  with  a  barium  salt.  Here  electrolysis  shows  that  sulphate- 
ion  is  absent  and  that  the  negative  ion  is  C2H5SO4~~. 

In  the  same  way  every  other  ionic  material  may  be  shown  to  be  a 
substance  with  an  individual  set  of  physical  (p.  375)  and  chemical 
properties.  Each  salt,  when  dissolved,  gives  two  kinds  (at  least)  of 
ionic  materials.  The  solution  is  simply  a  mixture,  and  each  physical 
component  forthwith  behaves,  towards  ions  capable  of  uniting  with 
it,  as  if  it  were  alone.  The  other  materials,  ionic  and  molecular, 
which  are  present,  may  remain  essentially  unaffected  throughout  the 
change. 

Application  in  Chemical  Analysis.  —  Since  the  larger  num- 
ber of  ordinary  chemical  substances  are  ionogens,  and  the  most  rapid 
and  simplest  chemical  changes  take  place  when  they  are  in  solution, 
the  various  reactions  of  their  solutions  are  employed  as  tests  for  the 
substances  in  question.  An  advantage  of  the  use  of  the  solutions  is 
that  they  contain  a  mixture  of  two  independent  materials,  the  anion 
and  the  cation,  and  when  these  have  been  identified  successfully  the 
salt  from  which  they  were  formed  is  known.  The  simplicity  to  which 
chemical  analysis  is  thus  reduced  may  be  seen  when  we  consider  that 
twenty-five  common  metals  with  twenty-five  negative  radicals  might 
give  a  total  of  over  six  hundred  different  salts.  If  the  distinct 
properties  of  each  of  these  had  to  be  considered,  the  identification  of 
an  unknown  substance  would  be  very  difficult.  In  solution,  how- 


386  INORGANIC  CHEMISTRY 

ever,  the  problem  becomes  much  easier.  Every  solution  made  from 
a  single  salt  will  contain  but  two  substances  (in  the  main;  see, 
however,  below),  and  the  problem  reduces  itself  to  ascertaining  which 
two,  out  of  a  total  of  fifty,  are  present  in  any  particular  case.  This 
is  easier  than  investigating  six  hundred  possibilities. 

As  an  example  of  the  method,  let  us  suppose  that  we  look  first 
for  the  positive  ion.  Most  systems  of  analysis  begin  by  the  addition 
of  a  solution  containing  chloride-ion,  generally  dilute  hydrochloric 
acid,  to  the  liquid.  If  an  ion  is  present  which  in  combination  with 
chloride-ion  gives  an  insoluble  compound,  a  precipitate  will  appear. 
Amongst  the  common  positive  ions  but  three  are  of  this  kind,  namely, 
silver-ion,  mercurous-ion,  and  lead-ion.  So  that  the  precipitate,  if 
it  appears,  is  a  chloride  of  one  of  these  three  metals,  and  the  matter 
of  distinguishing  between  the  three  is  quickly  disposed  of  by  further 
examination  of  its  properties.  If  no  precipitate  comes  out,  then  these 
three  metals  are  probably  absent,  and  some  fresh  ion  capable  of 
precipitating  another  set  of  positive  ions  is  introduced  (see  Chap. 
XXXVII).  Thus  by  a  process  of  elimination  we  quickly  find  out 
whether  any  metal  ion  is  present,  and,  if  so,  precisely  which  one  it  is. 

The  language  of  analysis  is  frequently  somewhat  loose.  Thus  we  speak  of 
the  addition  of  a  silver  salt  to  a  solution  as  being  a  "test  for  chlorine."  As  a 
matter  of  fact,  it  is  not  a  test  for  chlorine.  It  is  not  intended  as  a  test  for  free 
chlorine,  nor  will  it  show  the  presence  of  chlorine  in  many  states  of  combination. 
It  is  simply  a  test  for  ionic  chlorine  Cl~~,  and  cannot  give  us  information  in  regard  to 
the  presence  or  absence  of  any  other  form  of  the  element.  So  the  wet-way  tests 
for  "copper,"  "silver,"  etc.,  so  called,  are  tests  for  the  ionic  forms  of  these  ele- 
ments, and  not  for  the  presence  of  the  element  in  every  form.  Even  the  two 
kinds  of  copper  and  mercury  ions,  Cu*"4",  Cu+,  Kg*"1",  Hg"*~,  must  be  classed  as 
distinct  substances.  Thus,  the  last  is  precipitated  by  chloride-ion  while  the 
second  last  is  not,  mercuric  chloride  HgCl2  being  soluble. 

NEUTRALIZATION 

Neutralization.  —  When  80  per  cent  sulphuric  acid  is  poured 
upon  solid  potassium  hydroxide,  much  heat  is  developed  and  clouds  of 
steam  arise.  The  solid  product,  when  freed  from  the  rest  of  the 
water,  is  potassium  sulphate: 

2KOH  +  H2S04  *=»  2H20  T  +  K2S04. 

With  any  other  pair  consisting  of  an  acid  and  a  base,  a  similar  interac- 
tion occurs  (cf.  p.  213),  water  and  a  salt  being  produced. 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          387 

A  double  decomposition  between  ionogens  is  always  reversible 
(p.  380),  and  so  we  should  expect  that  in  dilute  solution  the  inter- 
action of  an  acid  and  a  base  would  be  incomplete.  We  find,  however, 
that  this  particular  sort  of  action  almost  always  goes  to  completion. 
This  kind  of  action  is  called  neutralization,  because  both  acid  and 
base  are  .completely  consumed,  and  hydrogen-ion  and  hydroxide-ion 
are  alike  impossible  of  detection  in  the  resulting  mixture.  The 
solution  is  neutral  to  litmus. 


(>  87*7  )  (01/7) 
(8%)HC1^C1-    +H+°(92%)    teractions^of  acids  and  bases 

(9%)  NaOH  ^  Na+  +  OH"  (91%)    Jl  sh?wn  in.  the  *"M5^ 
jj  i|  The  lomzation  of  the  hydro- 

xrni       TT  r\  chloric  acid  reaches  0.92  in  a 

JNav^l       ±120  .  11,. 

f  ^  i  Q  Of  \      ft  f\c\O7  \  decmormal  solution,  and  goes 

farther  when  the  acid  is  di- 

luted  with  the  water  of  another  solution.  That  of  the  sodium 
hydroxide  similarly  goes  beyond  0.91.  Thus  the  substances  in  the 
solutions  before  mixing  are  almost  entirely  ionic.  The  crosswise 
union,  H+  +  OH~~  *=?  H2O,  however,  is  all  but  complete,  for  water 
is  hardly  ionized  at  all  (p.  367).  The  materials  on  whose  inter- 
action with  the  Cl~  and  Na+,  respectively,  the  maintenance  of 
molecules  HC1  and  NaOH  depends,  being  thus  removed,  the  disso- 
ciation of  the  acid  and  base  promptly  brings  itself  to  completion, 
and  the  left  sides  of  the  equations  vanish.  Practically  all  the 
hydrogen-ion  and  hydroxide-ion  become  water,  which  thenceforth  is 
simply  a  part  of  the  solvent.  The  Cl~  and  Na+,  however,  if  the 
solution  is  now  1/20  normal,  unite  to  the  extent  of  0.13  only.  If 
it  is  more  dilute,  this  union  forms  a  still  smaller  factor  in  the  whole 
change.  Practically  it  is  negligible.  Now  all  that  has  been  ^said  of 
this  acid  and  base  will  apply  mutatis  mutandis  whenever  any  active, 
highly  ionized  acid  and  base  come  together.  Thus  we  may  write  one 
simple  equation  for  all  neutralizations  of  active  acids  and  bases: 

H+  +  OH"  ->  H20, 

without  omitting  anything  essential. 

The  ions  of  a  salt  are  always  left  over  from  the  main  action,  and 
may  be  brought  together,  in  turn,  by  evaporation:  Na+  +  Cl~  —  » 
NaCl,  or  the  liquid  may  be  used  as  a  solution  of  the  pure  salt. 

The  equations  as  commonly  writtc  Q  . 

NaOH  +  HC1  -»  NaCl  +  H2O, 
2NaOH  +  H2SO4 


388  INORGANIC  CHEMISTRY 

apply  to  the  interactions  when  water  is  absent.  If  used  for  neutralization  in 
dilute  solution,  it  must  be  understood  that  they  condense  two  changes  into  one 
equation.  The  formation  of  water  comes  first,  that  of  the  salt  afterwards. 
Sometimes  neutralization  is  wholly  misconstrued  by  the  supposition  being  made 
that  it  occurs  in  consequence  of  a  great  tendency  to  salt  formation. 

Confirmations  of  this  View  of  Neutralization.  —  The  neu- 
tral mixture  of  the  acid  and  base  gives  no  evidence  of  the  presence 
either  of  the  hydrogen  ions  or  of  the  hydroxide  ions.  The  character- 
istic tastes,  and  actions  upon  indicators,  of  these  two  ions,  and  the 
interaction  of  the  former  of  the  two  with  metals  like  magnesium, 
are  all  wanting.  That  this  is  due,  not  simply  to  two  opposing  in- 
fluences having  destroyed  each  other's  effects,  but  to  a  real  disappear- 
ance of  the  agencies  themselves,  may  be  demonstrated  by  showing 
that  the  total  number  of  ions  is  very  much  smaller  in  the  mixture 
than  in  the  two  substances  taken  separately. 

The  trough  (Fig.  106,  p.  363)  is  half-filled  with  a  dilute  solution 
(say,  N/4)  of  some  active  acid,  such  as  hydrochloric  acid.  An  equal 
volume  of  a  JV/4  solution  of  some  soluble  base  (loaded  with  sugar), 
such  as  sodium  hydroxide,  is  then  allowed  to  flow  in,  below  the  acid. 
On  completing  the  circuit  we  find  a  considerable  deflection  of  the 
amperemeter  (say,  1.5  amperes).  When  the  interaction  is  now 
brought  about  by  stirring,  a  very  great  fall  in  the  reading  (say  to 
0.5  amperes)  is  observed.*  The  only  plausible  explanation  is  that, 
not  only  have  many  of  the  ions  assumed  a  molecular  form,  but  those 
which  have  suffered  in  this  respect  have  been  the  most  rapidly 
moving  and  best  conducting  ones,  namely,  the  hydrogen-ion  and 
hydroxide-ion. 

Again,  a  considerable  thermal  effect  accompanies  neutralization. 
But,  in  the  cases  we  are  discussing,  that  is  where  active  bases  and 
acids  are  employed,  the  heat  liberated  by  use  of  equivalent  weights 
(p.  182)  is  always  the  same,  namely  13,700  cal.  That  it  is  always 
the  same  confirms  our  theory,  for  practically  the  whole  change  is 
always  the  formation  of  18  g.  of  water  from  the  ions. 

When  less  highly  ionized  acids  or  bases  are  used,  the  only  differ- 

*  The  experiment  may  be  made  more  striking  by  adding  a  few  drops  of 
phenolphthalem  solution  to  the  acid  and  using  a  minute  excess  of  the  base.  To 
prevent  the  appearance  of  a  pink  layer  at  the  interface,  and  before  the  stirring,  a 
thin  layer  of  sodium  chloride  solution  (loaded  with  less  sugar  than  the  solution 
of  the  base)  may  be  introduced  below  the  acid,  before  the  layer  of  the  base  is 
added 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          389 

ence  is  that  there  are  more  of  the  molecular  materials  present, 
before  the  solutions  are  mixed.  But  the  removal  of  the  H+  and 
OH~  ions  permits  the  molecules  of  the  acid  and  base  to  dissociate, 
so  that  the  final  products  are  water  and  the  ions  of  a  salt,  as  before. 
This  case  is  discussed  in  detail  later. 

The  foregoing  formulation  and  explanation  apply  to  every  case  of 
mixing  ionogens,  where  a  very  slightly  ionized  substance  is  one  of 
the  products,  that  is,  when  water,  or  a  feeble  acid,  or  a  feeble  base 
(pp.  369-370)  is  formed. 

Acidimetry  and  Alkalimetry.  When,  as  is  constantly  the 
case,  a  chemist  desires  to  ascertain  the  quantity  of  an  acid  or  base 

present  in  a  solution,  he  uses  for  the  purpose  \:he  interaction  just 
discussed.  If,  for  example,  the  problem  ic  to  "^certain  the  weight 
of  hydrogen  chloride  per  liter  in  a  specimen  of  hydrochloric  acid, 
this  can  be  done  by  neutralizing  a  measured  portion  of  this  acid  with 
a  solution  of  an  alkali  of  known  concentration  (see  next  section). 
The  volume  of  the  latter  which  is  required  for  the  purpose  is  ob- 
served. If  the  alkali  is  sodium  hydroxide,  the  action  taking  place  is 

HC1  +  NaOH  ->  H20  +  NaCl. 

The  volume  of  acid  is  measured  out  into  a  beaker  by  means  of  a 
pipette  (Fig.  107)  of  fixed  capacity,  which  is  filled  by  suction  to  the 
mark  on  the  stem.  Suppose  the  amount  to  be  25  c.c.  The  standard 
alkali  solution  is  placed  in  a  burette  (Fig.  108),  which  is  filled  down 
to  the  tip  of  the  nozzle.  A  few  drops  of  litmus  solution  are  now 
added  to  the  acid,  the  level  of  the  alkali  in  the  burette  is  read  off, 
and  the  alkali  is  allowed  to  run  slowly  into  the  acid.  After  a  time, 
the  hydroxide-ion  which  this  introduces  will  begin  to  produce  a  blue 
color,  close  to  where  the  stream  enters  the  liquid.  This  is  at  first 
dissipated  by  stirring,  and  the  whole  remains  red.  Finally,  however, 
a  point  is  reached  at  which  the  entire  solution  assumes  a  tint  inter- 
mediate between  blue  and  red.  With  one  drop  less  of  the  base,  it 
is  distinctly  red.  With  one  drop  more,  it  would  become  distinctly 
blue.  Litmus  paper  of  either  shade  dipped  in  this  neutral  solution 
remains  unaffected.  The  level  in  the  burette  is  read  again,  and 
the  difference  between  this  and  the  previous  reading  gives  the  number 
of  c.c.  of  standard  alkali  used. 

By  the  use  of  a  standard  solution  of  an  acid  in  the  burette,  the 
quantity  of  a  base  may  be  determined  in  the  same  way. 


390 


INORGANIC  CHEMISTRY 


FIG.  107. 


FIG.  108. 


Standard  Solutions.  —  The  standard  solutions  used  in  this 
work  are  usually  normal,  and  contain  one  equivalent  weight  of  the 
alkali  or  acid  in  one  liter  of  the  solution.  For  more  delicate  work, 
decinormal  (N/W)  solutions  may  be  employed.  The  concentra- 
tion of  such  a  solution  is  called  its  titer,  and  the  operation  of 
analyzing  another  solution  by  means  of  it,  titration.  The  value  of 
standard  solutions  lies  in  the  fact  that,  when  once  the  solution  has 
been  prepared,  and  the  exact  concentration  adjusted  by  quantita- 
tive experiments,  its  use  does  not  require  any  weighing,  and  the 
measurements  of  volumes  can  be  carried  out  with  great  rapidity. 
A  process  involving  weighing  need  not  again  be  undertaken  until  the 
stock  of  the  standard  solution  is  exhausted. 

The  calculation  of  the  result  is  also  simple.     One  liter  of  normal 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          391 

alkali  contains  17  g.  of  available  hydroxyl,  and  one  liter  of  normal 
acid,  1  g.  of  available  hydrogen  (p.  182).  Equal  volumes  of 
normal  solutions  will  therefore  exactly  neutralize  one  another,  18  g. 
of  water  being  formed  by  interaction  of  a  liter  of  each.  If,  for 
the  neutralization  of  the  25  c.c.  of  hydrochloric  acid  used  above, 
50  c.c.  of  normal  alkali  are  required,  the  acid  is  twice-normal  (2JV). 
When  15  c.c.  are  required,  the  acid  is  Jf  or  §]V.  If  the  actual 
weight  of  the  acid  in  the  latter  case  has  to  be  calculated,  we  remem- 
ber that  there  are  36.46  g.  of  hydrogen  chloride  in  1  1.  of  a  normal 
solution,  and  therefore  36.46  X  f  X  Tf tv  g-  =  0.5467  g.  in  25  c.c. 
of  a  solution  which  is  f-normal. 

Methods  of  quantitative  analysis  in  which  standard  solutions  are 
employed  are  known  as  volumetric  methods,  and  are  much  used  by 
analysts  and  investigators.  They  occupy  much  less  time  than 
gravimetric  operations,  in  which  numerous  weighings  have  to  be 
made,  and  are  often  just  as  accurate.  The  substances,  like  litmus, 
by  whose  change  of  color  the  completeness  of  the  action  is  made 
known,  are  called  indicators  (see  below). 

Indicators.  —  Indicators  are  substances  which,  in  presence  of 
certain  other  substances,  assume  a  very  deep  color,  or  change 
sharply  from  one  deep  color  to  another.  Thus,  phenolphthalein 
(p.  388)  is  colorless  in  presence  of  acids  (i.e.,  hydrogen-ion),  and  red 
(when  dilute,  pink)  in  presence  of  alkalies  (i.e.,  hydroxide-ion).  Lit- 
mus, again,  is  red  with  acids,  and  blue  with  alkalies.  The  change  of 
color  depends  upon  a  chemical  interaction  in  each  case,  but  since 
indicators  are  chosen  for  their  strong  coloration,  the  quantity  of 
the  acid  or  base  used  up  in  changing  the  tint  of  the  trace  of  the 
indicator  is  so  small  as  to  be  negligible.  The  common  indicators 
are: 

Phenolphthalein  C2oHi4O4,  a  colorless  substance  and  very  feeble 
acid.  It  is  not  perceptibly  dissociated  into  its  ions : 

C2oHi4O4  (colorless)  <=»  C2oHi304~  (red)  +  H+, 

and  in  neutral  or  acid  solutions  is,  therefore,  without  visible  color. 
When  a  base  is  added  gradually  to  an  acid  containing  some  of  this 
indicator,  the  acid  is  first  neutralized.  Then,  and  not  till  then,  the 
slightest  excess  of  hydroxide-ion  unites  with  the  trace  of  hydrogen- 
ion  from  the  phenolphthalein,  the  above  equilibrium  is  displaced 
forwards,  and  a  visible  amount  of  the  red  negative  ion  is  formed: 


392  INORGANIC  CHEMISTRY 

C20H1404  (colorless)  <=$  C20H13O4-  (red)*  +     H+  )  <_  „  n 
NaOH  ±9  Na+  +  OH"  J  ~ 

In  this  more  compact  formulation,  we  show  the  product  (H2O) 
from  the  union  of  the  two  ions  which  combine,  but  omit  the  prod- 
uct from  the  union  of  Na+  and  C20Hi3O4~,  because  here  (since  the 
product  is  a  salt)  hardly  any  union  occurs. 

This  indicator  is  especially  sensitive  to  acids  (weak  or  strong), f 
and  it  shows  the  presence  of  an  excess  of  alkali  most  sharply  when 
the  alkali  is  an  active  one  like  sodium  hydroxide,  and  should,  there- 
fore, be  employed  only  with  strong  bases.  With  a  weak  base  like 
ammonium  hydroxide,  a  considerable  excess  of  the  base  must  often 
be  used  before  the  color  appears. 

Litmus  is  an  extract  from  certain  lichens,  first  used  by  Boyle. 
It  contains  azolitmin.  One  of  its  colors  is  that  of  the  molecule, 
and  the  other  that  of  the  ion. 

Methyl  orange,  (CH3)2NC6H4.N :  N.C6H4SO3Na,  is  a  complex  or- 
ganic compound  which  gives,  in  acid  solution,  a  red  and  in  alkaline 
solution  a  yellow  color.  This  indicator  is  very  sensitive  to  bases, 
both  strong  and  weak.f 

Congo  red  is  the  sodium  salt  of  an  acid  of  complex  structure  (see 
Dyes).  In  neutral  or  alkaline  solutions  it  is  red;  with  acids  it  turns 
blue.  Paper  dipped  in  Congo  red  differs  from  litmus  paper  in  that  it 
shows  gradations  in  color,  the  blue  being  much  more  distinct  with  an 
active  acid  than  with  a  relatively  weak  one  like  acetic  acid  (p.  369). 
Litmus  paper  is  equally  red  with  all  acids  save  the  very  feeblest. 

Some  special  indicators  have  been  mentioned.  Thus,  starch 
emulsion  is  used  for  recognizing  the  presence  of  traces  of  iodine 
(p.  276).  Potassium  permanganate  is  itself  so  strongly  colored  that 
it  is  its  own  indicator  (p.  320). 

WEAK  ACIDS  AND  BASES 

Active  and  Weak  Acids.  —  In  solutions  containing  equivalent 
quantities  of  hydrogen  salts,  and  therefore  equal  amounts  of  com- 

*  The  ion  has  this  composition,  but,  in  reality,  has  a  different  chemical  struc- 
ture from  the  corresponding  part  of  the  original  molecule.  An  internal  rearrange- 
ment, not  representable  in  the  equation,  accompanies  the  dissociation.  The 
same  remark  applies  to  the  other  indicators. 

t  A  neutral  solution  contains  H+  and  OH~~  in  the  same  concentrations  as  in 
water,  namely  0.061AT.  Methyl  orange  becomes  yellow  with  0.041ATH+,  and 
phenolphthaleln  becomes  colorless  when  OH~  becomes  O.OsliV. 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          393 

bined  hydrogen,  in  equal  volumes,  the  concentration  of  hydrogen- 
ion  at  any  moment  in  each  will  be  different.  This  concentration 
will  be  high  or  low  according  to  the  extent  to  which  water  is  able 
to  dissociate  the  molecules.  Now  the  activity  of  the  hydrogen-ion, 
that  is,  the  speed  with  which  it  will  interact,  like  that  of  any  other 
substance,  depends  on  its  concentration  (p.  292).  Hence  the  hydro- 
gen salts  furnish,  on  being  dissolved,  acids  of  all  degrees  of  activity. 
Thus  in  normal  hydrochloric  acid,  the  fraction  dissociated  is  0.78, 
and  the  hydrogen-ion  is  0.78-normal,  whereas  in  normal  acetic  acid 
the  hydrogen-ion  is  only  0.004-normal  (p.  366).  Yet  the  amounts 
of  hydrogen  chloride  and  hydrogen  acetate  per  liter  contain  equal 
quantities  of  replaceable  hydrogen,  namely,  1  g.  each.  Both  the  solu- 
tions in  fact  are  normal  in  respect  to  combined  hydrogen.  But  the 
normal  acetic  acid  has  only  about  one  two-hundredth  of  the  activity 
of  normal  hydrochloric  acid. 

That  a  difference  in  the  activity  of  different  acids  does  exist  may 
be  shown,  roughly,  by  placing  similar  pieces  of  the  same  metal,  say 
zinc,  in  equal  volumes  of  various  normal  solutions  of  acids,  such  as 
hydrochloric,  sulphuric,  and  acetic.  The  hydrogen  is  evolved  more 
rapidly  by  the  first  than  by  the  second,  and  very  much  faster  by 
either  than  by  the  last.  Naturally,  the  first  is  sooner  exhausted, 
while  the  third  acts  in  its  slow  way  for  a  very  long  time  before  being  all 
used  up.  In  the  third  case  few  ions  of  hydrogen  are  at  hand  at  any 
one  moment,  but  more  are  formed  continuously  from  the  molecules,  to 
take  the  place  of  those  displaced.  Thus  the  total  amount  of  hydrogen 
obtained  from  each  acid  is  finally  the  same.  It  is  the  speed  of  evolu- 
tion alone  which  is  different  and  shows  the  differing  concentrations  of 
the  hydrogen-ion. 

In  cases  of  extremely  small  ionization,  the  presence  or  absence  of 
visible  action  on  litmus  may  form  another  means  of  estimating 
activity.  Thus,  litmus  is  easily  turned  red  by  a  deci-normal  solution 
of  acetic  acid  or  of  any  more  active  acid  (p.  369),  but  hydrogen  sul- 
phide, in  a  solution  of  the  same  molecular  concentration,  contains  only 
one-twentieth  as  many  hydrogen  ions  (p.  369),  and  affects  litmus 
paper  but  slightly.  Paper  dipped  in  Congo  red  exhibits  differences  in 
the  activity  of  acids  by  the  different  depths  of  the  tints  it  assumes. 
For  example,  it  is  much  less  markedly  affected  by  acetic  than  by 
sulphuric  acid  of  the  same  concentration  (Indicators,  p.  392). 

Many  hydrogen  salts  are  but  slightly  soluble.  Thus,  with  silicic 
acid  (q.v.),  the  solid  can  keep  only  a  small  concentration  of  molecules 
in  solution:  H2SiO3  (solid)  <=>  H<>Si03  (dslvd).  So  that,  although 


394  INORGANIC  CHEMISTRY 

some  ions  are  doubtless  present,  H2Si03  (dslvd)  <=*  2H+  +  Si03=,  their 
concentration,  being  dependent  on  that  of  the  molecules,  is  very  minute 
indeed.  Still,  even  in  the  absence  of  an  effect  upon  litmus,  the  sub- 
stance can  be  recognized  to  be  an  acid.  Thus,  by  the  action  of  sodium 
hydroxide,  silicic  acid  can  be  made  into  sodium  silicate  Na2SiO3, 
which  is  highly  soluble  and  highly  ionized.  Hence,  since  SiOs^  is  a 
negative  ion,  we  reach  the  conclusion  indirectly  that  H2SiO3  is  an 
acid. 

Substances  like  ammonia  NH3,  sugar,  and  alcohol,  although  they 
contain  hydrogen,  are  not  hydrogen  salts.  They  are  not  ionogens 
(c/.  p.  342),  and  give  no  hydrogen-ion.  lonizable  and  non-ionizable 
hydrogen  may  even  be  contained  in  the  same  compound.  Thus, 
each  molecule  of  acetic  acid  H.CO2CH3  contains  four  hydrogen  units, 
but  gives  only  one  hydrogen  ion.  The  other  three  are  part  of  the 
acetate-ion  CC>2CH3~.  We  infer  this  because  metals  can  be  substi- 
tuted for  one  hydrogen  unit  (NaCO2CH3),  but  not  more. 

Active  and  Weak  Bases.  —  The  case  of  weak  bases  is  exactly 
analogous  to  that  of  weak  acids.  Thus,  in  normal  potassium  hydrox- 
ide KOH,  the  fraction  ionized  is  0.77  and  the  hydroxide-ion  OH~ 
is  0.77-normal,  whereas  in  normal  ammonium  hydroxide  NH4OH 
the  hydroxide-ion  is  only  0.004-normal.  Yet  both  solutions  are 
normal  in  respect  to  ionized  and  combined  hydroxyl  together.  But 
normal  ammonium  hydroxide  has  only  about  one-one  hundred  and 
ninetieth  of  the  basic  activity  of  normal  potassium  hydroxide. 

Most  bases,  as  we  have  seen  (p.  374),  are  but  slightly  soluble,  and 
do  not  give  sufficient  hydroxide-ion  to  affect  indicators.  Yet  they 
interact  with  acids,  often  giving  soluble,  and  in  all  cases  highly 
ionized  salts. 

Trinitrobenzene  CeH3(N02)3,  which  is  colorless  in  acid  solution 
and  deep-orange  in  presence  of  an  active  base,  can  be  used  to  show 
different  concentrations  of  hydroxide-ion.  Thus,  with  JV-sodium 
hydroxide  it  gives  a  dark  orange  color  (0.737V  OH~),  but  with  a 
TV/ 10  solution  light  orange  (0.0917V  OH~).  With  TV-ammonium 
hydroxide  it  gives  a  very  faint  color  (0.004JV  OH~)  and  with  a 
N/IQ  solution  no  color  (0.0013N  OET). 

Neutralization  of  Little  Ionized  Substances.  —  When  con- 
centrated solutions  are  employed,  or,  when  acids  and  bases  which  are 
but  little  ionized  are  involved,  the  mechanism  of  the  change  is  still 
the  same  in  all  respects.  The  only  difference  is  that,  since  the  acid  or 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          395 

base  is  not  fully  ionized  to  start  with,  its  molecules  must  dissociate 
progressively,  in  proportion  as  the  hydrogen  ions  pass  into  combina- 
tion. All  the  hydrogen  and  hydroxyl  capable  of  forming  ions  will 
pass  through  that  stage,  and  ultimately  become  water,  before  the 
solution  can  reach  the  neutral  condition. 

The  neutralization  of  inactive  bases  and  acids  gives  instructive 
results,  when  performed  in  the  conductivity  trough  (Fig.  106,  p.  363). 
Thus,  if  the  trough  be  first  half-filled  with  AT/10  acetic  acid,  and  an 
equal  volume  of  N / 10  ammonium  hydroxide  (loaded  with  sugar) 
be  run  in  below  the  acid,  the  formulation  (two  horizontal  lines) 
shows  the  degrees  of  ionization : 

(>85%)         (0%) 

(98.7%)  HCO3CH3^  CO2CH3-  +  H+  (1.3%) 
(98.7%)NH4OH    i=>  NH4+         +  OH-  (1.3%) 

tl  tl 

NH4CO2CH3        H20 
«15%)          (100%) 

The  conductivity  before  mixing  (horizontal  lines)  is  very  slight. 
After  mixing,  water,  which  is  hardly  ionized  at  all,  and  ammonium 
acetate,  a  highly  ionized  salt  are  formed.  On  the  whole,  the  number 
of  ions  is  greatly  increased,  and  the  conductivity  increases  also  very 
markedly. 

From  all  this  it  will  be  seen  that  the  activity  of  acids  and  bases 
cannot  be  measured  by  the  quantity  of  base  or  acid  required  to  neu- 
tralize them.  The  full  amount  required  by  the  equation  is  always 
needed  in  every  case.  This  is  because  neutralization  uses  up  the 
hydrogen-  or  hydroxide-ion  at  once,  and  so  permits  the  rapid  generation 
of  a  fresh  supply.  The  concentration  of  one  of  these  ionic  materials 
can  only  be  measured  by  some  action  which  uses  it  up  slowly  or  not 
at  all,  so  that  ionic  double  decompositions  are  excluded.  In  the 
action  of  metals  on  acids  (p.  128),  and  in  determining  conductivity 
(p.  362),  the  consumption  of  the  ions  is  slow,  and  hence  the  measure- 
ment can  be  made  in  these  cases.  Actions  which  consume  no  ions 
at  all  are  also  known,  and  are  used  in  measuring  activity  (see  Carbo- 
hydrates and  esters). 

The  reader  must  not  fall  into  the  error  of  supposing  that  the 
neutralization  of  inactive  bases  and  acids  takes  a  longer  time  than 
that  of  active  ones.  The  formation  of  more  ions,  by  the  ionization 
of  the  molecules,  is  so  rapid  that  the  time  it  occupies  is  in  all  cases 
too  short  to  be  measured.  Ionization  appears  to  be  instantaneous. 


396  INORGANIC  CHEMISTRY 

When  the  acid  or  base  is  but  little  soluble  in  water,  as  when  zinc 
hydroxide  is  treated  with  a  dilute  acid,  one  other  link  is  added  to  the 
network  of  equilibria.  The  acid  proceeds  to  interact  with  the  small 
dissolved  part  of  the  base.  As  this  is  disposed  of,  solution  goes  on 
progressively  and,  through  a  train  of  equilibria : 

Zn(OH)2  (solid)  <=>  Zn(OH)2  (dslvd)  ?±  Zn++  +  20H~, 

the  supply  of  hydroxide-ion  is  maintained  until  all  the  molecules  of 
the  base,  solid  and  dissolved,  are  used  up  and  the  action  is  completed. 
Heating  hastens  these,  as  it  does  all  other  changes. 

If  acid  and  base  are  alike  insoluble,  it  saves  time,  if  the  production 
of  the  salt  is  the  ultimate  object,  to  fuse  the  materials  together  at  a 
high  temperature. 

Thermochemistry  of  Neutralisation.  —  The  above  interpre- 
tation of  the  phenomena  of  neutralization  is  confirmed  by  many  facts. 
Thus  a  considerable  amount  of  heat  is  liberated  in  neutralization. 
Now,  as  we  have  seen  (p.  388),  when  active  acids  (p.  373)  and  bases 
(p.  374)  in  dilute  solution  are  concerned,  it  is  found  that  the  quantities 
of  heat  for  the  neutralization  of  the  same  amount  of  hydrogen-ion,  or 
hydroxide-ion,  are  always  the  same,  namely,  13,700  cal.  for  equiv- 
alent weights.  If  the  action  consisted  primarily  in  the  formation 
of  a  different  salt  from  every  pair,  we  should  expect  the  heat  liberated 
to  be  different.  Thus,  the  heats  of  formation  of  dry  potassium 
chloride  and  dry  sodium  iodide  are  104,300  cal.  and  69,100  cal., 
respectively.  But  the  heats  of  formation  of  their  solutions  by 
neutralizing  the  proper  acids  and  bases  are  identical.  If,  however, 
in  such  cases,  neutralization  consists  always  simply  in  the  formation 
of  water,  we  should  expect  the  quantities  of  heat  liberated  to  be 
identical,  as,  in  fact,  they  are : 

H+  +  OH-  ->  H2O  +  13,700  cal. 

We  are  confirmed  in  these  conclusions  when  we  employ  concentrated  solu- 
tions, or  use  less  completely  ionized,  or  insoluble  acids  and  bases  for  neutraliza- 
tion. With  such  substances  —  and  they  are  in  the  majority  —  the  heats  of 
neutralization  are  not  alike,  but  different  in  every  case.  Thus,  for  dilute  solu- 
tions of  sodium  hydroxide  and  hydrofluoric  acid,  the  latter  a  slightly  ionized, 
soluble  acid,  the  thermochemical  equation  is  as  follows: 

NaOH  +  HF  -»  H2O  +  NaF  +  16,270  cal. 

Since  the  sodium  fluoride  is  fully  ionized,  the  only  difference  between  this  case  and 
the  preceding  one  is  that  the  hydrogen  fluoride  is  largely  in  the  molecular  condi- 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          397 

tion  to  start  with,  and  that  here,  in  addition  to  the  union  of  hydrogen  and 
hydroxyl  ions,  we  have  a  continuous  dissociation  of  the  hydrofluoric  acid  accom- 
panying the  neutralization.  The  fact  that  here  the  heat  produced  is  much  greater 
than  before,  shows  that  the  dissociation  of  this  acid  is  associated  with  the  produc- 
tion of  heat  (cf.  pp.  305,  368).  When  the  same  base  is  used  with  hypochlorous 
acid,  the  divergence  is  in  precisely  the  opposite  direction  and  about  the  same  in 
amount : 

NaOH  +  HC1O  -» H2O  +  NaCIO  +  9840  cal. 

Here  again  the  salt  produced,  sodium  hypochlorite,  is  fully  ionized,  so  that  the 
diminished  evolution  of  heat  must  be  due  to  the  fact  that  the  feebly  ionized 
hypochlorous  acid  (in  N / 10  solution,  only  0.0002  ionized)  absorbs  part  of  the 
heat  of  neutralization  in  passing  into  the  ionic  condition.  Applying  this  to  bases, 
we  find  that  the  neutralization  of  ammonium  hydroxide,  a  feebly  ionized  base, 
with  any  active  acid,  gives  a  heat  of  neutralization  below  the  normal: 

NH4OH  +  HC1  ->  H2O  +  NH4C1  +  12,200  cal. 

Here  again  the  salt  produced  is  fully  ionized.  Thus  the  ionization  of  the  ammo- 
nium hydroxide  must  have  consumed  an  appreciable  part  of  the  heat  of  neutraliza- 
tion which  would  otherwise  have  reached  the  normal  figure  of  13,700  calories. 

Volume  Change  in  Neutralisation.  —  When  the  volumes  of  the 
solutions  of  active  acids  and  alkalies  are  carefully  measured  before  being  mixed, 
and  compared  with  the  volume  of  the  neutral  mixture,  an  expansion  is  always 
found  to  have  occurred.  When  one  liter  of  a  normal  solution  of  each  substance  is 
taken  at  starting,  the  volume  of  the  mixture  is  always  20  c.c.  greater  than  that  of 
the  component  liquids.  This  is  rather  remarkable,  because  the  chemical  change 
has  produced  only  18  c.c.  of  water,  yet  the  volume  of  the  water  appears  to  be 
20  c.c.  greater  than  that  of  the  ions  from  which  it  was  formed.  This  shows  that 
the  change  involves  more  than  the  mere  union  of  the  ions.  The  electric  charges 
on  the  ions  cause  a  high  internal  pressure,  called  electrostriction,  especially  in 
the  molecules  of  water  immediately  surrounding  each  ion,  and  the  liquid  is  com- 
pressed. When  the  ions  unite,  the  pressure  is  removed,  and  expansion  occurs. 
When  less  highly  ionized  acids  and  bases  are  used,  the  alteration  in  volume  is 
irregular,  since  it  is  affected  by  the  occurrence  of  other  changes  than  the  mere 
union  of  hydrogen-ion  and  hydroxide-ion. 

Interaction  of  Salts  with  Acids  and  Bases.  —  When  a  highly 
ionized  acid  is  mixed  with  a  salt,  a  reversible  action  tending  to 
form  another  acid  and  salt  is  set  up.  Such  an  action  is  that  of  nitric 
acid  on  a  hypochlorite  in  dilute  solution : 

HNO3  +  KOC1  *=?  KN03  +  HOC1, 

giving  potassium  nitrate  and  hypochlorous  acid.  In  such  a  case,  if 
the  products  are  both  as  highly  ionized  as  the  initial  substances,  the 


398  INORGANIC  CHEMISTRY 

result  is  similar  to  that  of  the  interaction  between  potassium  chloride 
and  sodium  nitrate  (p.  381).     No  decisive  change  takes  place. 

With  hypochlorous  acid,  however,  which  is  very  slightly  ionized 
(0.02  per  cent  in  JV/10  solution),  the  result  is  different: 


ISSf^ocf  I  ^HOC1(dslvd)* 

This  acid  is  promptly  formed  from  its  ions,  and  the  final  mixture  con- 
tains, mainly,  K+,  NO3~  and  molecular  HOC1.  Yet,  since  the  sub- 
stance is  soluble,  no  outward  evidence  that  the  action  differs  from  that 
of  potassium  chloride  and  sodium  nitrate  is  visible.  The  conductivity, 
however,  is  found  to  have  been  greatly  reduced  when  the  solutions 
are  mixed  (p.  388),  because  half  the  ions,  including  the  most  rapidly 
migrating  of  the  four  (hydrogen-ion),  have  been  removed  (p.  347). 
The  action  of  an  active  acid  upon  a  solution  of  sodium  peroxide 
(p.  315)  is  another  illustration  of  this  sort  of  action. 

When  the  molecules  of  the  resulting  acid  are  insoluble,  then  it  may 
be  precipitated  (cf.  silicic  acid),  after  the  manner  of  silver  chloride 
(p.  382),  or  may  escape,  if  an  insoluble  gas  (cf.  hydrogen  sulphide), 
irrespective  of  its  degree  of  ionization. 

In  the  same  way,  when  a  salt  and  a  base  are  brought  together,  a 
base  and  a  salt  are  produced.  All  that  has  been  said  in  the  preceding 
paragraph  applies  to  this  case  also.  Thus  ammonium  hydroxide 
(q.v.),  which  is  a  feebly  ionized  base  (p.  370),  is  formed  on  this  plan, 
by  mixing  solutions  of  an  ammonium  salt  and  a  strong  base: 


NaOH  fc*  Na+  -f  +-  T^TT  nw  ^  i  A\ 


When  the  resulting  base  is  insoluble,  like  zinc  hydroxide,  it  is  pre- 
cipitated, and  the  action  becomes  nearly  complete  on  this  account  and 
irrespective  of  the  degree  of  ionization. 

Interaction  of  Salts  with  Water:  Hydrolysis.  —  The  natural 
ionization  of  water  is  very  slight,  but  there  are  cases  in  which  its 
effects  become  noticeable,  and  the  interaction  of  its  ions  with  those 
of  dissolved  salts  cannot  be  neglected.  For  example,  an  aqueous 

*  To  save  space,  this  mode  of  formulation  (p.  392)  will  be  used  when  there  is 
very  complete  union  of  only  one  pair  of  the  ions.  Here  the  K+  and  NO3~,  being 
the  ions  of  a  salt,  combine  to  a  very  slight  extent  only. 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          399 

solution  of  pure  cupric  sulphate  is  always  acid  and  therefore 
contains  excess  of  hydrogen-ion: 

CuS04  *=»  S04=  +  Cu++  )  <__  r   ,nTrx    ,A  ,    ,>. 
2H20  ±*  2H+  +  20H-  |  *  Cu(OH)2  (dslvd)* 

Cupric  hydroxide,  being  a  very  feeble  base,  and  comparable  with 
water  itself  in  the  small  extent  to  which  the  solvent  is  able  to  hold  its 
ions  apart,  is  formed  to  a  small  extent.  The  removal  of  some  hydrox- 
ide-ion by  this  means  enables  more  of  the  water  to  dissociate.  This 
in  turn  furnishes  the  material  for  the  production  of  more  cupric 
hydroxide.  The  action  does  not  proceed  very  far,  but  it  makes 
sufficient  progress  to  leave  a  perceptible  excess  of  hydrogen-ion  in 
the  liquid  and  to  give  it,  therefore,  an  acid  reaction.  The  hydrogen- 
ion  combines  slightly,  but  only  slightly,  with  the  sulphate-ion,  for 
sulphuric  acid  is  a  highly  ionized  acid.  This  part  of  the  action  has, 
therefore,  been  left  out  of  the  diagram.  The  ordinary  equation  for 
this  change  would  be: 

CuS04  +  H20  <=+  Cu(OH)2  +  H2S04. 

The  hydrolysis  is  much  greater  with  sodium  sulphide  (q.v.)  and  anti- 
mony trichloride  (q.v.). 

Again,  soap  solution  is  always  faintly  alkaline: 

Na+  +  C15H31C02 
OH-  +       H+ 

The  sodium  palmitate  is  highly  ionized,  but  palmitic  acid 
is  hardly  ionized  at  all.  The  final  result  is  the  production  of  a  recog- 
nizable amount  of  hydroxide-ion  in  the  solution.  Thus,  a  salt  derived 
from  an  acid  and  a  base  of  very  different  degrees  of  activity,  whether 
it  is  the  base  (as  Cu(OH)2)  or  the  acid  (as  palmitic  acid  or  hydrogen 
sulphide,  q.v.)  which  is  the  weaker  member,  is  likely  to  be  more  or 
less  hydrolyzed  by  water.  In  the  former  case  the  solution  is  acid,  in 
the  latter  basic  in  reaction.  When  both  the  acid  and  the  base  are 
weak,  the  hydrolysis  is  more  extensive.  Other  things  being  equal, 
salts  containing  bivalent  or  trivalent  radicals  are  more  noticeably 
hydrolyze.d  than  are  those  composed  only  of  univalent  radicals. 

It  will  be  seen  that  hydrolysis  is  the  precise  reverse  of  neutraliza- 
tion (pp.  389,  395).  The  latter  being  almost  always  nearly  complete, 
the  former  must  be,  as  a  rule,  very  slight. 

Cases  of  this  kind  being  common,  we  are  thus  compelled  to  enlarge 
our  list  of  possible  components  in  the  solution  of  any  salt  (cf.  375). 


£00  INORGANIC  CHEMISTRY 

In  addition  to  the  molecules  and  ions  of  the  salt,  there  are  present, 
water  and  its  ions,  and  the  molecules  of  the  base  and  acid  formed  by 
the  union  of  the  latter  ions  with  the  former.  There  are  thus  no  less 
than  eight  different  components  in  the  mixture. 

MlXED    lONOGENS   AND   DOUBLE   SALTS 

As  a  rule,  a  univalent  ion,  such  as  chlorate-ion  C103~~,  unites  with 
one  kind  of  cation  to  give  but  one  kind  of  salt.  The  result  is  called 
a  neutral  or  normal  salt,  as  KC1O3  or  NaClO3.  The  acid,  chloric 
acid,  is  called  a  monobasic  acid,  for  its  molecule  reacts  with  but  one 
molecule  of  a  base.  The  possibilities  are  more  numerous,  however, 
with  an  ion  of  higher  valence.  Thus: 

GARBONATE-ION  MAY  GIVE:  CALCIUM-ION  MAY  GIVE: 

H2CO3,  the  acid,  Ca(OH)2,  the  base, 

Na2C03,  a  neutral  salt,  CaCl2,  a  neutral  salt, 

NaHC03,  an  acid  salt,  Ca(OH)Cl,  a  basic  salt,* 

NaKC03,  a  mixed  salt,  CaCl(OCl),  a  mixed  salt. 

Carbonic  acid  is  a  di-basic  acid,  and  calcium  hydroxide  a  di-acid  base. 
The  last  two  compounds  of  each  set  are  mixed  ionogens.  Their  char- 
acteristic is  that  they  contain  more  than  two  kinds  of  radicals  and 
break  up  in  solution,  giving  more  than  two  kinds  of  ions. 

Acid  Salts.  —  The  acid  salts  may  be  obtained  by  using  half 
that  quantity  of  the  base  which  would  be  required  fully  to  neutralize 
the  acid,  and  evaporating  the  resulting  solution: 

NaOH  +  H2S04  ^  H20t,+  NaHS04. 

With  a  monobasic  acid,  say  hydrochloric  acid,  this  treatment  gives 
simply  a  mixture  of  the  normal  salt  and  the  free  acid,  and  not  a  single 
substance. 

Acid  salts  are  also  formed  by  the  interaction  of  other  salts  with  an 
excess  of  the  acid  (pp.  206,  207). 

The  acid  salt  is  intermediate  in  composition  between  the  acid  itself 
and  the  normal  salt.  All  of  the  hydrogen  of  the  acid  has  not  been 
displaced  by  the  metal.  It  is  named  an  acid  salt  on  account  of  its 
composition,  but  is  not  necessarily  acid  in  its  reaction  towards 
litmus.  That  depends  on  whether  its  solution  contains  a  sufficient 
amount  of  hydrogen-ion  to  affect  indicators.  Sodium-hydrogen 
sulphate  gives  the  ions  Na+  and  HSO4~,  but,  even  in  moderately  dilute 

*  This  particular  basic  salt  has  not  been  isolated  in  a  pure  state. 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          401 

solution,  the  latter  ion  is  further  dissociated  into  H+  and  SCU"  to  a 
large  extent  (p.  367).  Its  solution  is  therefore  acid  in  truth.  On  the 
other  hand,  sodium-hydrogen  carbonate  NaHCO3,  derived  from 
carbonic  acid,  H^COs,  gives  the  ions  Na+  and  HCOs",  and  the  amount 
of  hydrogen-ion  formed  by  the  latter  is  too  small  to  be  detected  by 
indicators.  This  acid  salt  gives  therefore  a  solution  which  is  actually 
neutral  to  litmus.* 

Basic  Salts.  —  Corresponding  to  the  acid  salts  we  have  also 
basic  salts,  about  which  statements  parallel  to  the  above  might  be 
made.  Thus,  from  sodium  hydroxide  but  one  salt  can  be  formed. 
With  lead  hydroxide,  Pb(OH)2,  however,  the  displacement  of  one 
hydroxyl  by  a  negative  radical,  without  the  disturbance  of  the  other, 
is  conceivable  and  can  be  achieved.  The  half-chloride,  for  example, 
is  called  lead  oxy chloride  Pb(OH)Cl.  The  basic  salts  are  usually 
insoluble  in  water,  and  therefore  as  a  rule  do  not  exhibit  the  basic 
reaction  with  litmus. 

Mixed  Salts.  —  So-called  mixed  salts,  like  sodium-potassium 
carbonate  KNaCOs  (see  Silicates),  may  be  obtained  by  half  neutraliz- 
ing the  acid  with  the  calculated  amount  of  one  base  and  then  complet- 
ing the  operation  with  the  other.  Corresponding  treatment  will  give 
mixed  salts  of  a  di-acid  base. 

It  will  be  seen  that  2KNaC03  is  equivalent  to  K2C03,Na2C03,  and 
that  2CaCl(OCl)  is  equivalent  to  CaCl2,Ca(OCl)2.  Since  we  have  as 
yet  no  general  means  of  determining  the  molecular  weights  of  solids, 
there  is  no  generally  applicable  way  of  deciding  which  formula  is 
preferable  (see,  however,  Bleaching  powder,  p.  475,  and  under 
Calcium).  In  solutions  of  these  salts  the  ions  which  are  found  might 
come  from  a  substance  possessing  either  of  the  alternative  formulae, 
so  that  no  light  is  thrown  on  the  question  by  this  means.  Thus, 
most  compounds  of  this  kind,  with  the  exception  of  acid  and  basic 
salts,f  are  considered  to  be  molecular  compounds  (p.  154)  of  two 
salts  and  are  classed  as  double  salts. 

*  Because  of  hydrolysis  (p.  398),  the  solution  of  the  "neutral"  salt,  sodium 
carbonate  Na2COs,  is  actually  alkaline  in  reaction.  The  terms  "acid,"  "basic," 
and  "neutral,"  applied  to  salts,  refer  simply  to  the  composition  and  ignore  the 
behavior. 

t  The  formulae  of  basic  salts  even  are  often  written  as  if  such  salts  were 
molecular  compounds,  as  Cu(OH)2,CuCl2,  or  even  CuO,CuCl2,H2O,  in  place  of 
Cu(OH)Cl  (see  Copper). 


402  INORGANIC  CHEMISTRY 

Double  Salts.  —  Substances  similar  to  ferrous-ammonium  sul- 
phate FeS04,(NH4)2S04,6H20,  and  alum  (q.v.),  are  very  numerous. 
Because  their  formulae  can  be  written  so  as  to  show  two  complete 
salts,  and  because  they  are  easily  formed  by  crystallization  from  a 
solution  containing  both  salts,  they  are  called  double  salts.  In 
solution  they  are  resolved  into  their  constituent  salts,  and  these,  in 
turn,  are  ionized.  Almost  always  the  acid  radicals  are  identical 
(see,  however,  Kainite). 

Each  kind  of  ion  of  a  double  salt  exhibits  its  own  properties,  irre- 
spective of  the  nature  of  the  numerous  substances,  ionic  and  otherwise, 
which  are  present.  Hence,  when  a  solution  of  a  particular  ionic 
material  is  required,  solutions  of  such  bodies  are  often  used  instead  of 
those  of  simpler  ones,  if  for  any  reason  the  substitution  is  convenient. 
The  choice  of  the  complex  compound  must  be  made  in  such  a  way 
that  the  other  ions  shall  not  interfere  with  that  particular  reaction  of 
one  of  them  which  is  in  question. 

The  class  of  bodies  known  as  salts  of  complex  adds  (q.v.)  are  ion- 
ized like  ordinary  salts  and  not  like  double  salts. 

KINDS  OF  IONIC  CHEMICAL  CHANGE 

Five  distinct  varieties  of  chemical  change  are  characteristic  of 
ionic  materials.  These  are:  (1)  Disunion  or  combination  of  ions, 
(2)  displacement  of  the  material  of  one  ion  by  another  substance,  (3) 
destruction  or  formation  of  a  compound  ionic  material,  (4)  change  in 
the  charges  of  two  ionic  materials,  (5)  charge  or  discharge  of  two  ionic 
materials,  in  electrolysis.  Every  one  of  these  kinds  of  action  has  been 
illustrated,  some  of  them  very  frequently,  in  the  present  and  foregoing 
chapters. 

1.  Disunion  and  Combination  of  Ions.  —  This  sort  of  change 
is  illustrated  every  time  an  ionogen  is  dissolved  in  water  (disunion)  or 
a  solution  of  such  a  substance  is  evaporated  (combination).  Both  of 
the  directions  of  this  sort  of  change  occur  also  to  a  greater  or  less 
extent  whenever  solutions  of  two  ionogens  are  mixed.  In  the  latter 
case: 

(1)  Two  salts  give  two  salts  (pp.  380-382). 

(2)  An  acid  and  salt  give  a  salt  and  an  acid  (p.  397). 

(3)  A  base  and  salt  give  a  salt  and  a  base  (p.  398). 

(4)  An  acid  and  base  give  water  and  a  salt  (neutralization). 

(1)  is  complete  only  when  at  least  one  product  is  insoluble.  (2) 
and  (3)  are  complete  when  at  least  one  product  is  little  ionized  or 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          403 

insoluble  or  both.  (4)  is  almost  always  complete  because  water  is 
generally  less  ionized  than  any  other  substance  in  the  system. 

2.  Displacement  of  One  Ion:  Electromotive  Series  of  the 
Metals.  — The  displacement  of  hydrogen  from  dilute  acids  by 
metals  (p.  118)  now  appears  in  a  new  light.  These  interactions  do 
not  occur  in  the  absence  of  water  (p.  119).  The  action  will  be  formu- 
lated thus: 

Zn  +  2H+  +  S04=  -r*  Zn++  +  SO4=  +  H2. 

The  sulphate-ion  S04=,  although  zinc  sulphate  is  somewhat  less  ion- 
ized than  sulphuric  acid,  is  not  much  affected  by  the  change  and  may 
be  omitted: 

Zn  +  2H+  -+  Zn++  +  H2. 

Thus,  this  action,  which  takes  place  in  the  same  fashion  with  most 
acids,  concerns  in  reality  only  the  hydrogen-ion  in  the  solution,  and 
is  independent  of  the  nature  of  the  negative  ion.  True,  hydrogen-ion 
cannot  be  used  alone,  for  it  is  always  accompanied  by  some  negative 
radical.  But  the  latter,  like  the  vessel  in  which  the  experiment  is 
made,  is  part  of  the  necessary  apparatus,  and  not  an  interacting  sub- 
stance. The  action  consists  simply  in  the  transference  of  the  electric 
charges  from  the  hydrogen  to  the  zinc,  whereby  the  latter  becomes 
ionic.  The  discharged  hydrogen  is  liberated  as  gas.  The  speed  of 
the  action,  for  any  one  metal,  depends  on  the  concentration  of  the 
hydrogen-ion,  i.e.,  on  the  activity  of  the  acid.  Hence  the  action  goes 
faster  with  hydrochloric  acid  than  with  acetic  acid.  When  the 
solution  is  evaporated,  the  ionogen,  in  the  above  case  zinc  sulphate, 
is  formed : 

Zn++-f-SOr=-+ZnS04. 

Similarly,  fluorine  displaces  chlorine  from  chloride-ion  (p.  281), 
chlorine  displaces  bromine  from  bromide-ion  (p.  273) : 

C12  +  2Br~  -»  2Cr  +  Br2, 

and  bromine  displaces  iodine  from  iodide-ion  (p.  279).  Each  of  these 
actions  is  independent  of  the  nature  of  the  other  ion  which  accom- 
panies the  one  undergoing  change. 

The  same  sort  of  displacement  occurs  with  all  positive  ions.  Thus, 
zinc  will  not  only  displace  hydrogen,  but  also  other  metallic  elements, 
like  lead,  copper,  and  silver,  from  the  ionic  condition  in  solutions  of 
their  salts: 

Zn  +  Cu++  -+  Zn++  +  Cu  J  . 


404 


INORGANIC  CHEMISTRY 


Here  the  copper  appears  as  a  red  precipitate.  Lead,  in  turn, 
will  displace  copper  and  silver,  but  not  zinc.  Copper  will  displace 
silver.  Thus  the  metals  can  be  set  down  in  an  order  such  that  each 
metal  displaces  those  following  it  in  the  list  and  is  displaced  by  those 
preceding  it.  This  list  (see  below)  is  known  as  the  electromotive 
series  of  the  metals,  because  in  electrolysis  of  normal  solutions  of 
their  salts,  the  electromotive  force  of  the  current  required  to  deposit 
each  metal  (cf.  p.  359)  is  less  than  that  for  the 
metal  preceding  in  the  list  (see  Chap.  XXXVIII). 
This  list  embodies  many  facts  in  the  behavior 
of  the  metals,  and  should  be  kept  in  mind  as 
furnishing  a  key  to  the  actions  in  which  a  free 
metal  is  used  or  produced.  For  example,  the 
chemical  activity  (p.  129)  of  the  free  metals 
places  them  in  the  same  order.  The  earliest 
ones  rust  much  more  readily  in  air  than  do  the 
later  ones,  provided  the  oxide  does  not  adhere  so 
firmly  to  the  surface  of  the  metal  (e.g.,  Al)  as  to 
protect  it  from  further  action  of  the  oxygen  in 
the  air.  Those  following  copper  do  not  rust. 
Conversely,  the  oxides  of  the  metals  down  to 
and  including  manganese,  when  heated  in  a 
stream  of  hydrogen,  may  give  lower  oxides,  but 
are  not  completely  reduced  to  the  metallic  con- 
dition. The  oxides  of  cadmium  and  succeeding 
metals  are  easily  reduced.  The  oxides  of  mer- 
cury and  the  last  four  metals  are  decomposed 
by  heating  alone.  The  relations  of  the  metals 
in  respect  to  combination  with  elements  other 
than  oxygen  are  similarly  expressed  by  the  ar- 
rangement in  this  table. 

The  position  of  hydrogen  is  particularly  sig- 
nificant. It  will  be  noted  that  none  of  the  metals 
preceding  hydrogen  are  found  free  in  nature  as 
ordinary  minerals,*  while  all  of  the  metals  succeeding  hydrogen, 
although  occurring  to  some  extent  in  combination,  are  found  also 
free.  The  explanation  of  this  is  that,  by  prolonged  action  upon 
ordinary  water,  containing,  as  it  must,  carbonic  acid  and  other 
sources  of  hydrogen  ions,  the  metals  preceding  hydrogen  must 

*  Free  lead  and  tin  do  occur  as  rare  minerals.     Iron,  with  a  little  cobalt  and 
nickel,  constitutes  many  meteoric  masses. 


ELECTROMOTIVE 

SERIES  OF  THE 

METALS. 

Alkali  metals  (q.v.) 
Alkaline  earth 
metals  (q.v.) 
Magnesium 
Aluminium 
Manganese 
Zinc 

Chromium 
Cadmium 
Iron 
Cobalt 
Nickel 
Tin 
Lead 

Hydrogen 
Copper 
Arsenic 
Bismuth 
Antimony 
Mercury 
Silver 
Palladium 
Platinum 
Gold 


IONIC  SUBSTANCES   AND  TflElR  INTERACTIONS          405 

'eventually  displace  the  hydrogen-ion  and  pass  into  some  form  of 
combination  (cf.  p.  403).  The  metals  following  hydrogen  do  not 
displace  hydrogen-ion  and  are  much  less  affected  by  the  agencies 
which  are  most  active  in  the  chemical  transformation  of  minerals, 
Hence  they  often  remain  in  the  free  state.  For  this  reason  gold, 
silver,  and  copper  were  the  metals  first  used  by  man.  Iron  came 
into  service  much  later. 

To  avoid  a  common  misconception,  it  must  be  noted  that  the  electromotive 
series  has  no  bearing  on  the  tendency  of  one  radical  to  dislodge  another  in  double 
decompositions.  The  place  of  an  element  in  the  E.M.  series  defines  its  relative 
activity  when  free,  and  has  to  do  only  with  actions  where  one  free  element  dis- 
places (p.  129)  another.  The  influences  which  determine  a  double  decomposi- 
tion (cf.  pp.  208,  382-384)  are  such  as  the  insolubility  of  a  compound,  and  have 
little  to  do  with  the  chemical  activity  of  the  compounds  concerned  (p.  304)  and 
nothing  whatever  to  do  with  that  of  the  free  elements,  for  these,  in  fact,  are  not 
present  at  all.  Thus,  when  silver  chloride  (insoluble)  is  placed  in  a  solution  of 
potassium  iodide  KI,  it  is  quickly  converted  into  silver  iodide  Agl.  This  hap- 
pens because  silver  iodide  is  still  more  insoluble  than  is  silver  chloride.  But  free 
chlorine  (the  gas)  will  quickly  displace  the  iodine  from  silver  iodide. 

The  negative  ions  can  be  arranged  in  order  in  a  similar  way. 

3.   Destruction  or  Formation  of  a  Compound  Ion.  —  The 

destruction  of  a  compound  ionic  material  is  observed  in  the  action  of 
any  reducing  agent,  such  as  hydrogen  peroxide  (p.  320) ,  upon  a  dilute 
solution  of  a  permanganate.  The  compound  ion  MnC^"  gives  by 
reduction  Mn++  and  water.  It  also  occurs  when  charcoal  is  added 
to  hydrogen  peroxide  solution  (p.  318),  for  the  ion  O2=  of  the  peroxide 
gives  the  ion  O=  (or  OH~)  of  water  and  free  oxygen : 

4H+  +  202=  -»  2H2O  +  02. 

The  converse  occurs  when  potassium  cyanide  is  added  in  excess  to 
a  solution  of  a  salt  of  silver.  First,  silver  cyanide  is  precipitated,  and 
then  this  compound  unites  with  the  excess  of  cyanide-ion : 

Ag+  +  N03~  +  K+  +  NC"  ^  K+  +  N03~  +  AgNC  J, , 
K+  +  NO"  +  AgNC  ?->  K+  +  Ag(NC)2- 

The  product  is  the  soluble  potassium  argenticyanide.  It  is  a  salt  of 
the  complex  acid  HAg(CN)2,  and  not  a  double  salt  (p.  402).  It  does 
not  decompose  into  potassium  and  silver  cyanides  and  their  ions  when 
in  solution,  for  the  second  action,  above,  is  not  appreciably  reversible. 
Actions  of  this  class  often  proceed  slowly,  so  that  their  speed  can 
be  measured. 


406  INORGANIC  CHEMISTRY 

4.  Change  in  the   Charges  of  Two  Ions.  —  A  decrease  in 
the  charge  in  two  ions  probably  occurs  in  the  preparation  of  chlorine 
(p.  219).     The  decomposition  of  the  manganese  tetrachloride  takes 
place  by  the  simultaneous  discharge  of  two  equivalents  of  electricity 
from  the  quadrivalent  manganese  ion  and  two  chloride  ions: 

Mn++++  +  4C1-  _»  Mn++  +  2C1-  +  C12. 

Both  this  sort  of  change  and  its  converse  are  common  with  ions  of 
metals  such  as  iron  and  tin  (q.v.),  which  have  more  than  one  valence. 

5.  Charge  and  Discharge  of  Two  Ions,  Electrically.  —  Dis- 
charge of  ions  is  brought  about  in  every  electrolysis  (p.  216).     Thus, 
when  hydrochloric  acid  is  decomposed  by  the  current,  we  have: 

2H+  +  20  ->  H2      and      2Q-  +  20  -» C12. 

The  converse  takes  place  when  the  polarization  current  (p.  360)  is 
allowed  to  flow.  Both  charge  and  discharge  occur  in  every  simple 
battery,  as  when  zinc  dissolves  in  dilute  sulphuric  acid  to  give  zinc 
sulphate  (pp.  29-30) : 

Zn  4  Zn++  +  20       and      2H+  ->  H2  +  20 . 

The  creation  of  the  positive  charges  on  the  zinc-ion  in  the  former 
of  the  two  equations  leaves  the  rod  of  zinc  negatively  charged. 
The  liberation  of  the  positive  charges  from  the  hydrogen  ions  in  the 
latter  renders  the  platinum  wire  positive. 

NON-IONIC  MODES  OF  FORMING  IONOGENS 

While  ionogens  may  always  be  made  by  the  union  of  the  proper 
ions,  they  must  nevertheless,  in  the  absence  of  the  solvent,  be  regarded 
as  chemical  substances  which  may  be  constructed,  and  very  frequently 
are  made,  out  of  their  constituents  without  reference  to  the  ionic  plane 
of  cleavage.  Thus  we  have  incidentally  observed  many  ways  in 
which  acids,  bases,  and  salts  may  be  prepared  that  do  not  involve  a 
union  of  the  constituent  ions  and  are  not  ionic. 

Acids  and  Bases.  —  Oxygen  acids  can  almost  all  be  prepared 
from  the  anhydride,  that  is,  from  the  oxide  of  the  non-metal,  which 
is  not  an  ionogen,  and  water.  Phosphoric  acid,  sulphurous  acid 
(p.  149),  and  many  other  acids  are  so  formed.  The  hydrogen  halides 
are  all  producible  by  union  of  the  constituent  elements.  Some 
acids  are  formed  from  others  when  the  latter  are  exposed  to  light, 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS         407 

as  hydrochloric  acid  from  hypochlorous  acid  (p.  223);  or  when 
heated,  as  are  chloric  acid  from  hypochlorous  acid  (q.v.),  and 
phosphoric  from  phosphorous  acid  (q.v.).  Bases  are  formed  by 
the  union  of  oxides  of  metals  with  water  (p.  149). 

Salts.  —  The  dry  ways  of  forming  salts  are  very  numerous. 
Thus,  many  are  formed  by  direct  union  of  the  elements,  as  in  the 
case  of  chlorides  (p.  221),  sulphides  (p.  16),  and  other  simple  salts. 
Many  are  made  by  reduction  or  oxidation  from  other  salts,  as 
potassium  chloride  from  potassium  chlorate  (p.  83),  or  potassium 
perchlorate  (q.v.)  from  the  latter.  Often  a  reducing  or  an  oxidizing 
agent  is  used,  as  in  making  sodium  nitrite  (q.v.)  from  the  nitrate,  and 
lead  sulphate  from  lead  sulphide  (q.v.),  no  solvent  being  present. 
Almost  all  oxygen  salts  can  be  obtained  by  the  union  of  two  oxides, 
as  calcium  carbonate  (q.v.)  from  calcium  oxide  and  carbon  dioxide. 
Ammonium  salts  are  formed  by  combination  of  ammonia,  which  is 
not  an  ionogen,  with  anhydrous  acids  (p.  212). 

In  manufacturing  salts,  non-ionic  methods,  like  the  above,  as  well 
as  those  involving  ionic  actions,  are  very  commonly  used.  In  each 
case  the  cheapest  and  most  easily  accessible  materials  are  chosen, 
and  the  least  expensive  operation  is  selected. 

Neutralization  is  theoretically  the  simplest  ionic  way  of  getting  a  salt,  because 
the  water  can  be  removed  by  mere  evaporation.  Yet  most  of  the  salts  which 
are  on  the  market  are  made  by  the  use  of  other  actions.  In  fact,  the  pure  bases 
and  acids  are  usually  too  expensive  to  be  utilized  as  sources  of  salts. 

The  commonest  definition  of  a  salt,  as  a  substance  formed  by  the  neutraliza- 
tion of  an  acid  by  a  base,  is  open  to  many  objections.  It  is  logically  defective 
because  it  does  not  describe  what  a  salt  is,  but  one  method  of  making  a  salt, 
which  is  an  entirely  different  matter.  It  is  unfortunate  in  its  choice  amongst 
possible  paralogisms,  because  neutralization  is  more  significant  as  a  method  of 
forming  water  than  as  a  means  of  preparing  a  salt.  And  finally,  as  we  have  just 
seen,  the  definition  has  not  even  the  excuse  of  practical  value,  for  most  salts  are 
manufactured  by  entirely  different  reactions. 

Exercises.  —  1.  Using  the  data  in  regard  to  ionization  (p.  367), 
formulate  other  dissociations  according  to  the  models  on  p.  377. 

2.  In  the  case  of  the  green  solution  of  cupric  bromide  (p.  379), 
explain  in  detail  (p.  359)  the  effect  of  the  addition  of  potassium 
bromide.     Formulate  the  action  (p.  380). 

3.  Give  a  list  of  all  the  colorless  ionic  substances  you  can  think  of 
(p.  375). 


408  INORGANIC  CHEMISTRY 

4.  Formulate  fully,  according  to  the  diagram  on  p.  382,  the 
precipitation  of  barium  sulphate  (p.  385),  of  silicic  acid  from  sodium 
silicate  (p.  394),  of  zinc  hydroxide  from  zinc  sulphate  (p.  396),  of 
silver  chloride  from  silver  sulphate,  and  the  liberation  of  hydrogen 
chloride  by  phosphoric  acid  (p.  207). 

5.  Give  a  list  of  the  specific  physical  and  chemical  properties 
(p.  384)  of  iodide-ion,  sulphate-ion,  cupric-ion,  chloride-ion. 

6.  Formulate  (p.  403)  the  displacement  of  iodine  by  bromine 
(p.  279),  and  of  bromine  by  chlorine  (p.  273). 

7.  Explain  the  acid  reaction  of  ferric  chloride  FeCl3  solution 
(p.  399). 

8.  Name  all  the  physical  components  in  aqueous  solutions  of  po- 
tassium hydroxide,  hydrogen  chloride,  and  sulphuric  acid  (cf.  p.  400). 

9.  Name  the  anions  and  cations  whose  formulas  are  used  on 
p.  375. 

10.  Formulate  (p.  403)  the  actions  of  iron  and  of  aluminium  on 
dilute  hydrochloric  acid. 

11.  What  is  the  molar  concentration  (p.  183)  of  hydrogen-ion 
in  AT/10  hydrogen  sulphide  (p.  367)  and  in  AT/10  acetic  acid,  of  sodium- 
ion  in  AT/2  sodium  chloride,  and  of  cupric-ion  in  N  cupric  nitrate? 

12.  Combining  the  models  on  pp.  396  and  387,  formulate  the 
action  of  hydrochloric  acid  on  magnesium  hydroxide  and  on  zinc 
hydroxide. 

13.  Formulate  (p.  381)  and  discuss  the  action  of  sulphuric  acid 
upon  potassium  permanganate  (p.  320). 

14.  Formulate  (p.  387)  the  neutralizations  on  pp.  396-397. 

15.  What  do  we  infer  (p.  401)  from  the  fact  that  the  solution  of 
sodium  hydrogen  sulphide  NaHS  is  neutral? 

16.  Invent  an  interaction  of  two  soluble  salts  in  which  both 
products  shall  be  insoluble  (see  Table  of  Solubilities,  inside  front 
cover). 

17.  To  which  classes  of  ionic  actions  do  those  of  iodine  on  hy- 
drogen sulphide  (p.  278),  and  of  magnesium  on  cold  water  (p.  115), 
belong?     Formulate  the  former  according  to  the  model  on  p.  403. 

18.  What  metals,  beside  platinum,  would  be  most  likely  to  form 
suitable  electrodes  for  an  electrolytic  cell  (p.  404)? 

19.  How  should  you  attempt  to  obtain  (p.  405)  a  pure  aqueous 
solution  of  the  acid  HAg(CN)2? 

20.  Formulate  (p.  406)  the  electrolysis  of  hydriodic  acid  and  that 
of  cupric  sulphate,  the  latter  between  copper  electrodes  (p.  360). 

21.  Give,  for  each  of  the  following,  two  definitions,  one  in  terms  of 


IONIC  SUBSTANCES  AND  THEIR  INTERACTIONS          409 

experimental  facts,  the  other  in  terms  of  the  ions:  acid,  base,  salt, 
neutralization,  acid  salt,  mixed  salt. 

22.  Using  the  table  of  fractions  ionized  (p.  367),  prepare  lists  of 
the  pairs  of  ionic  substances  which  show  the  greatest,  and  the  least 
tendency  to  combine,  and  state  in  each  case  the  proportion  com- 
bining in  decinormal  solution. 

23.  In  the  case  of  the  chocolate-brown,  concentrated  solution  of 
cupric  bromide  (p.  378),  explain  in  detail  what  would  happen  to 
the  system:    (a)  if  metallic  zinc  were  to  be  added  (p.  404);    (b)  if 
hydrogen  sulphide  gas  were  to  be  led  into  the  solution    (CuS    is 
insoluble).     Formulate  each  case. 

24.  What  is  implied  by  the  statements,  that  peroxides  are  salts 
and  that  hydrogen  peroxide  is  feebly  acid  (p.  393)? 

25.  Formulate  after  the  model  on  p.  398,  and  discuss  fully,  the 
interaction  of:    (a)  sodium  peroxide  and  hydrochloric  acid  (p.  315); 
(b)  barium  peroxide  and  sulphuric  acid. 

26.  For  the  neutralization  of  77  c.c.  of  a  certain  alkaline  solution, 
25  c.c.   of  normal  hydrochloric  acid  are  required.      What  is  the 
normal    concentration    of  the   alkali?      If  the   alkali  was  sodium 
hydroxide,  what  weight  of  the  substance  was  present?     If  the  alkali 
was  barium  hydroxide,  what  weight  of  it  was  present? 

27.  What   would    be   the    behavior  in  a    conductivity   trough 
(Fig.  106)  experiment,  in  which  the  layers  were:    (a)  hydrochloric 
acid  and  sodium  acetate  in  equivalent  amounts,  before  and  after 
mixing;    (b)   hydrochloric    acid   and  sodium  peroxide?     Formulate 
both  (p.  387)  actions. 


CHAPTER  XX 
SULPHUR  AND  HYDROGEN  SULPHIDE 

Occurrence.  —  Free  sulphur  is  found  in  volcanic  regions  in 
Sicily,  where  it  is  mixed  with  gypsum  and  other  minerals  and  occupies 
the  pores  of  pumice-stone.  Rocky  materials  accompanying  a  mineral 
in  this  way  are  called  the  matrix.  The  other  important  deposits  are 
in  Louisiana  and  in  Texas  (Brazonia  county).  There  are  many 
minerals,  compounds  containing  sulphur,  which  are  chiefly  important 
on  account  of  their  other  constituents.  Sulphides  of  metals,  such  as 
pyrite  FeS2,  copper  pyrites  CuFeS2,  galena  PbS,  zinc-blende  ZnS,  and 
sulphates,  like  gypsum  CaSO4,2H20,  barite  BaSC>4,  and  celestite 
SrSO4,  are  fairly  plentiful.  The  deposits  of  free  sulphur  are  believed 
to  have  been  formed  mainly  from  gypsum,  by  reduction.  Sulphur 
is  a  constituent  of  the  proteins,  which  are  important  components  of 
the  structures  of  plants  and  of  animals. 

Manufacture.  —  In  Sicily  sulphur  is  obtained  by  the  simple 
process  of  jnelting  it  away  from  the  accompanying  volcanic  rock  at  a 
low  temperature.  .  A  part  of  the  sulphur  itself  is  burned,  to  obtain 
heat  to  melt  the  rest,  because  no  coal  is  found  in  Italy,  and  fuels  other 
than  sulphur  are  locally  too  expensive.  The  liquid  sulphur  is  allowed 
to  run  into  wooden  molds,  in  which  it  solidifies  in  the  form  of  roll 
sulphur,  or  roll  brimstone.  For  many  purposes  this  sulphur  is  suffi- 
ciently pure.  To  produce  the  best  quality  it  is  subjected  to  distilla- 
tion from  earthenware  retorts.  The  vapor  passes  into  a  large  brick 
chamber  and  condenses  upon  the  walls  and  floor  in  the  form  of  a  fine 
powder,  sold  as  flowers  of  sulphur.  When  the  chamber  has  become 
heated,  the  sulphur  condenses  in  the  form  of  a  liquid,  which  is  drawn 
off  and  cast  in  molds  as  before. 

In  Louisiana,  the  sulphur  forms  a  deposit  over  half  a  mile  in  diam- 
eter, below  900  feet  of  clay,  quicksand,  and  rock.  It  is  extracted  by 
the  Frasch  method,  by  means  of  borings  which  permit  four  pipes,  one 
within  the  other,  to  reach  the  deposit.  Water,  previously  heated, 
under  pressure  to  170°,  is  pumped  down  the  two  outside  pipes  (6  and 
8  inches  in  diameter).  After  time  has  been  allowed  for  the  melting 

410 


SULPHUR  AND  HYDROGEN  SULPHIDE  41V 

of  a  mass  of  the  sulphur  (m.-p.  114.5°),  compressed  air  is  forced  down 
the  innermost,  one-inch  pipe.  The  melted  sulphur  has  twice  the 
specific  gravity  of  the  water  in  the  outer  pipes.  But  the  mixture  of 
air  and  sulphur  has  about  the  same  specific  gravity,  and  so  flows 
freely  up  the  three-inch  pipe  surrounding  the  air  pipe.  The.  element 
flows  into  a  large,  wooden  enclosure,  in  which  it  solidifies,  and  is 
practically  pure  sulphur.  Each  well,  until  obstructed  by  collapse  of 
the  rock  and  quicksand  at  the  bottom,  produces  500  tons  a  day. 

The  greater  part  of  the  sulphur  of  commerce  formerly  came  from 
Sicily,  where,  in  1898,  447,000  tons  were  manufactured  against  41,000 
tons  elsewhere.  The  whole  supply  of  the  United  States  (250,000 
tons)  is  now  obtained  from  Louisiana  and  Texas.  In  1913,  Sicily 
produced  407,307  tons  and  Japan  58,452  tons,  and  in  1914  an  island 
off  the  coast  of  New  Zealand  furnished  12,000  tons.  The  world's 
consumption  is  over  800,000  tons. 

Some  sulphur  is  also  obtained  from  the  exhausted  material  used  in 
removing  sulphur  from  illuminating-gas  during  its  purification. 

Physical  Properties.  —  The  chief  physical  peculiarity  of  sul- 
phur is  that,  instead  of  existing  in  three  familiar  physical  states  only, 
like  water,  it  possesses  two  familiar  and  perfectly  distinct  solid  forms 
and  two  different  liquid  states  of  aggregation. 

1.  Rhombic  Sulphur.     Native  sulphur  is  yellow,  has  a  sp.  gr.  2.06 
and  melts  at  112.8°.     It  is  almost  insoluble  in  water,  but  dissolves 
freely  in  carbon  disulphide  (41  parts  in  100  at  18°)  and  in  sulphur 
monochloride  (q.v.).     The  crystals  of  native  sulphur,  as  also  those 
obtained  by  evaporating  a  solution,  belong  to  the  rhombic  system 
(Fig.  8,  p.  14).     Roll  sulphur,  and  most  specimens  of  flowers  of  sul- 
phur, are  the  same  substance,  although  the  crystals  in  their  growth 
have  interfered  with  one  another,  and  the  mass  is  crystalline,  simply, 
and  not  well  crystallized.     This  variety  is  called,  from 

its  form,  rhombic  sulphur.  This  form  is  stable  below 
96°.  Above  that  temperature  it  changes  slowly  into 
monoclinic  sulphur. 

2.  Monoclinic  Sulphur.     When  a  large  mass  of 
melted   sulphur   solidifies   slowly,   and  the  crust  is 
pierced  and  the  remaining  liquid  poured  out  before 

the  whole  has  become  solid,  the  interior  is  found  to         FIQ  1Qg 
be  lined  with  long,  transparent  needles  (Fig.  109). 
This  kind  of  sulphur  is  nearly  colorless,  has  a  sp.  gr.  1.96,  melts  at 
119.25°,  and  is  in  all  physical  respects  a  different  individual  from 


412  INORGANIC  CHEMISTRY 

rhombic  sulphur.  This  variety  is  named,  from  the  system  to  which 
its  crystals  belong,  monoclinic  sulphur. 

Monoclinic  sulphur  can  be  kept  only  above  96°  and  below  its 
melting-point  (119.25°).  Every  recently  solidified  mass  of  sulphur 
is  composed  of  it.  But,  below  96°  the  mass  gradually  becomes 
opaque,  the  change  usually  spreading  from  the  points  at  which  it  has 
been  touched  and  finally  affecting  the  whole  mass.  The  opacity  is 
due  to  the  fact  that  the  material  has  turned  into  an  aggregate  of  small 
particles  of  rhombic  sulphur,  each  of  which  occupies  less  space  than 
the  monoclinic  sulphur  from  which  it  was  formed.  Conversely; 
rhombic  sulphur,  when  heatedabove  96°,  but  not  as  high  as  its  melting- 
point,  turns  slowly  into  monoclinic  sulphur.  Contact  with  a  piece 
of  monoclinic  sulphur,  or  mere  rubbing  with  a  hard  body,  will  deter- 
mine the  point  at  which  the  transformation  shall  begin,  and  the 
expansion  which  accompanies  this  results  in  a  spreading  opacity  as 
before.  The  delay  before  the  change  starts  and  the  effect  of  rubbing 
and  inoculation  are  familiar  in  connection  with  almost  all  changes  of 
state  (cf.  p.  194). 

Transitions,  marked  by  definite  points,  like  this  one  at  96°,  are 
attended  by  similar  phenomena,  whether  they  lie  between  two  solid 
states,  or  a  liquid  and  a  solid  state  (ice  and  water),  or  a  gaseous  and  a 
liquid  state  (steam  and  water) .  Heat  is  given  out  when  we  pass  in  one 
direction  and  absorbed  by  passage  in  the  other.  The  rate  of  change 
of  vapor  pressure  with  change  in  temperature  (cf.  p.  200)  is  different 
on  each  side  of  the  transition  point.  A  body  which  has  two  solid 
states,  and,  therefore,  two  crystalline  forms,  is  said  to  be  dimorphous 
(two-formed),  and  one  with  more  than  two  such  states  polymorphous 
(see  Ammonium  nitrate).  But  this  term  is  not  intended  to  imply 
that  the  relation  of  two  solid  states  to  each  other  is  essentially 
different  from  that  of  two  states  of  different  kinds,  such  as  solid 
and  liquid,  although  the  term  "polymorphous"  is  not  applied  to 
the  latter. 

3.  S\  and  SM;  Vapor.  When  melted  sulphur  is  heated,  it  under- 
goes a  gradual  change,  which  is  especially  noticeable  near  160°. 
The  formerly  pale-yellow,  mobile  liquid  (S\)  suddenly  becomes 
dark-brown  in  color  (color  chiefly  due  to  organic  impurities)  and  so 
viscous  (SM)  that  the  vessel  may  be  inverted  without  loss  of  material : 
Sx  ^  S^.  The  liquid  is  a  mixture,  containing  increasing  proportions 
of  SM,  namely,  at  120°  3.6  per  cent,  at  160°  11  per  cent,  and  at  444.7° 
over  30  per  cent.  Beyond  260°  the  viscidity  becomes  less,  and  at 
444.7°  the  liquid  boils  and  passes  into  sulphur  vapor. 


SULPHUR  AND  HYDROGEN  SULPHIDE  413 

Insoluble,  Amorphous  Sulphur  SM.  —  When  sulphur  which 
has  been  exposed  to  the  air  is  boiled  and  then  allowed  slowly  to  cool, 
the  product  is  crystalline  and  soluble  in  carbon  disulphide,  as  before- 
But  when  such  impure  sulphur  is  boiled  and  then  suddenly  chilled  by 
pouring  into  cold  water,  it  is  at  first  semi-fluid.  After  several  days 
this  plastic  sulphur,  as  it  is  called,  becomes  hard.  It  is  then  found  to 
contain  rhombic  sulphur  mixed  with  about  34  per  cent  of  another 
variety  of  free  sulphur,  namely,  SM.  This  variety  is  almost  insoluble 
in  any  solvent,  and  so  may  be  secured  by  washing  the  mixture  with 
carbon  disulphide. 

This  insoluble  sulphur,  being  without  crystalline  structure,  is 
called  also  amorphous  (Gk.  d  priv.,  pop<j>rj  form)  sulphur.  Now 
amorphous  bodies  (see  Glass)  are  always  supercooled  liquids,  that  is, 
liquids  still  existing  as  such  at  a  temperature  at  which  the  solid,  crys- 
talline form  is  the  stable  one.  They  have  been  brought,  by  cooling,  so 
rapidly  through  their  freezing-point,  that  crystallization  has  not  had 
time  to  begin  (cf.  p.  175)  and  a  general  rigidity  only  has  supervened. 
Now  amorphous  sulphur  is  viscous,  liquid  sulphur  SM  which,  by  sudden 
chilling,  has  been  carried  past  both  the  gradual  transition  to  mobile 
liquid  sulphur  S\,  and  the  crystallization  as  well,  without  under- 
going either  of  these  changes.  It  is  supercooled  SM.  This  accounts 
for  the  fact  that  it  is  obtainable  only  by  rapid  cooling.  Once  the 
mixture  has  been  obtained  by  chilling,  the  insoluble  sulphur  reverts 
very  slowly  to  the  soluble  variety,  and  years  are  required  for  the  com- 
pletion of  the  reversion  at  room  temperature.  At  100°  the  reversion 
is  completed  in  an  hour. 

The  capacity  of  the  SM  to  be  supercooled  at  all  seems  to  depend  on 
the  presence  of  traces  of  foreign  bodies.  Of  these,  iodine  is  the  most 
efficient.  The  sulphuric  acid,  produced  by  prolonged  exposure  of  the 
sulphur  to  the  air  (see  below) ,  is  the  agent  usually  responsible  for  the 
supercooling.  Freshly  recrystallized  sulphur  gives  no  plastic  sulphur 
and  no  insoluble  sulphur.  Destruction  of  the  trace  of  free  acid,  by 
leading  ammonia  gas 'through  the  boiling  sulphur,  likewise  gives  a 
liquid  which,  when  chilled,  yields  nothing  but  crystalline  sulphur  [Lect. 
exp.].  Iodine  and  sulphuric  acid  are  negative  catalysts. 

Insoluble  sulphur  is  sometimes  found  in  flowers  of  sulphur  and 
always  in  sulphur  formed  by  precipitation  from  thiosulphates  (q.v.) 
in  presence  of  acids. 

Chemical  Properties.  —  When  the  density  of  sulphur  vapor  is 
determined  at  low  temperatures  and  under  reduced  pressures,  the 


414  INORGANIC  CHEMISTRY 

molecular  weight  corresponds  closely  to  the  formula  83.  As  the  tem- 
perature is  raised,  however,  the  vapor  expands  very  rapidly,  and  at 
800°  the  molecular  weight  is  64.2,  and  the  formula  therefore  82. 
Intermediately,  mixtures  of  Ss,  Se,  and  S2  exist  (Preuner  and  Schupp) . 
The  formula  of  dissolved  sulphur,  as  measured  by  freezing-  and 
boiling-point  methods  (p.  335),  is  83. 

We  do  not  ordinarily  think  of  sulphur  as  a  very  active  chemical 
substance,  but  this  is  largely  due  to  the  fact  that  its  solid  condition 
interferes  with  the  attainment  of  close  contact  with  the  body  upon 
which  it  acts.  It  is  really  very  active.  When  finely  divided  metals, 
with  the  exception  of  gold  and  platinum  (cf.  p.  404),  are  rubbed 
together  with  powdered  sulphur,  union  takes  place  and  sulphides  are 
produced.  Sulphur,  when  heated,  combines  with  great  vigor  with 
iron  (p.  16),  copper,  and  most  of  the  metals.  Sulphur  unites  also 
with  many  of  the  non-metals.  Thus  with  oxygen  it  produces  sulphur 
dioxide  (p.  88),  and  even  sulphur  trioxide  863.  It  unites  also  with 
chlorine  directly.  When  sulphur  is  treated  with  oxidizing  agents  in 
presence  of  water,  no  trace  of  sulphur  dioxide  (or  sulphurous  acid) 
is  formed;  the  only  product  is  sulphuric  acid  (see  p.  446).  Even  the 
oxygen  in  the  air  slowly  oxidizes  sulphur,  with  the  help  of  atmospheric 
moisture,  to  give  sulphuric  acid,  28  +  2H20  +  3O2  — »  2H2SO4.* 

Uses  of  Sulphur.  —  Large  quantities  of  crude  sulphur  are  em- 
ployed for  making  sulphur  dioxide,  which  is  used  in  the  manufacture 
of  sulphuric  acid,  in  bleaching  feathers,  straw,  and  wool,  in  preparing 
dried  fruits,  and  in  making  alkali  sulphites  for  employment  in  the 
bleaching  industry  and  in  paper-making.  The  manufacture  of  carbon 
disulphide  consumes  much  sulphur.  Purified  sulphur  is  employed  in 
the  manufacture  of  gunpowder,  fireworks,  matches  and,  by  combina- 
tion with  rubber,  of  vulcanite.  Flowers  of  sulphur  is  used  in  vineyards 
to  destroy  fungi,  which  it  does  by  virtue  of  the  traces  of  sulphuric  acid 
it  yields  by  oxidation. 

HYDROGEN  SULPHIDE  H2S. 

This  compound  is  found  in  some  mineral  waters,  which  in  conse- 
quence are  known  as  sulphur  waters.  It  is  produced  in  the  decompo- 
sition of  animal  matter  containing  sulphur  (proteins),  when  air  is 
excluded,  and  the  odor  of  rotten  eggs  is  due  in  part  to  its  presence. 

*  The  paragraph  on  the  chemical  relations  of  the  element  (see  end  of  this 
ehapter)  should  be  read  at  this  point. 


SULPHUR  AND  HYDROGEN  SULPHIDE  415 

Preparation.  —  1.  Hydrogen  and  sulphur  do  not  unite  percep- 
tibly in  the  cold.  At  310°  almost  complete  union  occurs,  but  about 
168  hours  (7  days)  are  required  for  the  change. 

2.  Sulphides  of  metals,  being  salts,  are  acted  upon  more  or  less 
easily  by  dilute  acids  (p.  397),  and  give  hydrogen  sulphide.     Ferrous 
sulphide,  the  least  expensive  of  those  easily  affected,  is  generally  used: 

FeS  +  2HC1  -»  H2S  +  FeCl2. 

For  hydrochloric  acid  we  may  substitute  an  aqueous  solution  of  any 
active,  non-oxidizing  acid  (see  p.  416).  A  Kipp's  apparatus  (p.  119)  is 
commonly  employed.  The  theory  of  this  action  is  discussed  later 
(see  p.  419). 

3.  Hydrogen  sulphide  is  the  invariable  product  of  the  extreme  re- 
duction of  any  sulphur  compound.     Thus,  it  is  formed  by  the  action 
of  hydrogen  iodide  upon  concentrated  sulphuric  acid  (p.  277).     Even 
sulphur  itself  is  reduced  by  dry,  gaseous  hydrogen  iodide: 


The  action  appears  to  be  just  the  reverse  of  that  which  takes  place  in 
aqueous  solution  (p.  278),  but  in  reality  is  quite  different.  Iodine  and 
gaseous  hydrogen  sulphide  will  not  produce  free  sulphur  and  gaseous 
hydrogen  iodide,  for  this  action  would  involve  a  considerable  increase 
in  energy  in  the  system.  But,  in  water,  they  do  give  hydrogen-ion 
and  iodide-ion,  for  these  bodies  contain  very  much  less  energy  than 
does  hydrogen  iodide: 

2H+  +  S=  +  I2-^2H+  +  S  i  +  2I~      or      S=  +  I2->S  i  +  2I~. 


This  action  is  simply  an  ionic  displacement. 

Physical  Properties.  —  Hydrogen  sulphide  is  a  colorless  gas 
with  a  characteristic  odor.  When  liquefied,  it  boils  at  —  62°,  and  in 
solid  form  melts  at  —83°.  At  12°  the  liquid  exerts  a  pressure  of 
15  atmospheres.  The  solubility  in  water  at  10°  is  360  volumes  in  100, 
and  becomes  less  as  the  temperature  is  raised.  The  gas  can  be  driven 
out  completely  by  boiling  the  solution  (cf.  p.  211).  The  gas  is  very 
poisonous,  one  part  in  two  hundred  being  fatal  to  mammals,  and  more 
than  once  fatal  accidents  have  occurred  in  chemical  laboratories. 

Chemical  Properties  of  the  Gas.  —  When  heated,  the  gas  dis- 
sociates, and  is  therefore  not  very  stable: 


416  INORGANIC  CHEMISTRY 

At  310°  the  decomposition,  although  very  slow,  affects  a  small  but 
perceptible  proportion  of  the  gas  before  coming  to  rest.  The  dissoci- 
ation, like  most  thermal  dissociations,  is  accompanied  by  an  absorp- 
tion of  heat  and  is  therefore  greater  at  higher  temperatures  (cf.  p.  305)  . 
The  gas  burns  in  air,  forming  steam  and  sulphur  dioxide.  The 
temperature  of  the  mantle  of  flame  surrounding  the  gas,  as  it  issues 
from  a  jet,  being  far  above  310°,  the  gas  in  the  interior  is  dissociated 
before  it  meets  with  any  oxygen.  Hence  a  cold  dish  held  across  the 
flame  (Fig.  110)  receives  a  deposit  of  free  sul- 
phur, and  a  part  of  the  hydrogen  is  also  liable 
to  escape  unburnt.  It  may  be  remarked  that 
dissociation  of  this  kind  probably  precedes  the 
combustion  of  most  gaseous  compounds  (see 
Flame). 

The  metals,  down  to  and  including  silver  in 
the  electromotive  series,  when  exposed  to  the 
gas,  quickly  receive  a  coating  of  sulphide.     The 
FIG.  no.  tarnishing  of  silver  in  the  household  is  probably 

due  to  a  trace  of  hydrogen  sulphide  in  the  il- 
luminating gas  which  escapes  from  slight  leaks  in  the  pipes.  That 
the  gas  should  thus  behave  like  free  sulphur  shows  its  instability. 

This  instability  is  shown  also  in  the  fact  that  its  hydrogen  reduces 
substances,  such  as  sulphur  dioxide,  which  are  not  affected  by  free 
hydrogen  : 


This  action  takes  place  much  more  rapidly  when  the  gases  are  moist 
than  when  they  are  dry  (p.  97),  and  is  retarded  by  dilution  with 
indifferent  gases  (cf.  p.  291).  Native  sulphur  is  occasionally  pro- 
duced by  this  action  (see,  however,  p.  410,  par.  1),  as  both  of  these 
gases  are  found  issuing  from  the  ground  in  volcanic  neighborhoods. 
Sulphur  is  deposited  also  when  hydrogen  sulphide  undergoes  a  partial 
combustion  with  a  restricted  supply  of  oxygen:  2H2S  +  O2  —  > 
2H2O  +  2S. 

When  hydrogen  sulphide  gas  is  led  through  concentrated,  or  even, 
simply,  normal  sulphuric  acid,  the  acid  is  reduced,  sulphur  dioxide 
escapes,  and  sulphur  is  deposited  : 

H2S  +  H2S04  ->  S  +  2H20  +  S02. 

The  sulphuric  acid  may  be  written  H20,SO3.  In  furnishing  S02, 
therefore,  each  molecule  can  give  one  unit  of  oxygen  and  therefore 


SULPHUR  AND  HYDROGEN  SULPHIDE  417 

oxidize  one  molecule  of  H2S  (see  p.  426).  On  account  of  this  action 
the  gas  cannot  be  dried  by  means  of  concentrated  sulphuric  acid. 
Calcium  chloride  is  likewise  inapplicable,  since  a  partial  interchange 
takes  place,  resulting  in  the  production  of  calcium  sulphide  and 
hydrogen  chloride  gas.  Only  a  dehydrating  agent,  such  as  phos- 
phoric anhydride,  with  which  it  cannot  interact,  is  suitable  for  drying 
the  gas. 

A  Characteristic  of  Reduction  and  Oxidation.  —  In  the  first 
of  the  three  actions  last  mentioned,  it  will  be  seen  that,  while  the 
SO2  was  reduced  to  S,  at  the  same  time  H2S  was  oxidized  (to  S).  In 
the  second  action,  H2S  was  oxidized  to  S,  and  02  was  reduced  to 
2H2O.  In  the  last  action,  H2S  was  oxidized  to  S,  and  H2SO4  was 
reduced  to  SO2.  It  is  a  characteristic  of  such  actions  that,  one  sub- 
stance is  oxidized  and  another  reduced:  oxidation  and  reduction 
always  occur  together,  in  the  same  reaction.  Here,  under  hydrogen 
sulphide,  W3  speak  of  its  reducing  effect  on  sulphur  dioxide,  or  on 
sulphuric  acid.  Under  sulphur  dioxide  and  sulphuric  acid,  however, 
we  should  speak  of  the  oxidizing  effect  of  the  substance  on  hydrogen 
sulphide. 

Chemical  Properties  of  the  Aqueous  Solution  of  Hydrogen 
Sulphide.  —  While  the  gas  itself  is  not  an  acid,  its  solution  in 
water  gives  a  feeble  acid  reaction  with  litmus,  and  is  sometimes 
named  hydrosulphuric  acid  H2S,  Aq.  The  conductivity  of  a  2V/10 
aqueous  solution  is  small,  and  only  0.0007  (0.07  per  cent)  of  the 
substance  is  ionized: 

H2S  <=?  H+  +  HS-  (±=>  H+  +  S=). 

Some  S~  ions  are  present.  But  hydrosulphide-ion  HS~,  although 
an  acid,  is  less  dissociated  than  is  water  itself,  and  the  amount  of 
sulphide-ion  is  therefore  very  small.  The  salts  of  hydrosulphide- 
ion,  such  as  NaHS  (sodium  acid  sulphide,  see  next  section),  give 
therefore  neutral  solutions.  Thig  behavior  is  the  rule  with  the 
acid  salts  of  feeble  dibasic  acids  (p.  401). 

As  an  acid,  the  solution  of  hydrogen  sulphide  may  be  neutralized 
by  bases.  For  the  same  reason  it  enters  into  double  decomposition 
with  salts  (see  next  section). 

By  the  action  of  oxygen  from  the  air  upon  an  aqueous  solution  of 
hydrogen  sulphide,  the  sulphur  is  slowly  displaced  and  appears  in  the 
form  of  a  fine  white  powder : 

02  +  2H2S-*2S|  +2H20. 


418  INORGANIC  CHEMISTRY 

This  is  an  action  similar  to  the  displacement  of  ionic  iodine  by  free 
chlorine  (p.  279).  On  the  other  hand,  the  hydrogen  may  be  displaced 
by  metals,  particularly  the  more  active  ones,  but  the  small  degree  of 
ionization  makes  the  action  very  slow. 

The  solution  of  the  gas  is  a  reducing  agent,  as  its  action  upon 
iodine  shows  (p.  278).  So,  also,  in  presence  of  an  acid,  it  removes 
oxygen  from  dichromic  acid,  produced  by  the  action  of  an  acid  upon 
potassium  dichromate: 

K2Cr207  +  2HC1  <=±  H2Cr207  +  2KC1  (1) 

H2Cr2O7  +  6HC1  ->  4H20  +  2CrCl3  (+  30)  (2) 

(30)       +  3H2S  ->  3H20  +  3S (3) 

Adding:  K2Cr2O7  +  8HC1  +  3H2S-»2KC1  +  2CrCl3  +  7H2O+3S 

The  first  partial  equation  (cf.  p.  269)  represents  the  regular  interaction 
of  two  ionogens,  but  the  second  interaction  does  not  take  place  unless 
an  oxidizable  body  (here  the  hydrogen  sulphide)  is  present  to  take 
possession  of  the  oxygen  which  it  is  capable  of  delivering  (cf.  p.  320) . 
This  action  illustrates  the  decomposition  of  a  compound  ion  (p.  405). 
Here  Cr2O7=  gives  chromic-ion  Cr+++  and  water. 

Sulphides. —  As  a  di-basic  acid  (p.  372),  hydrogen  sulphide 
gives  both  acid  and  normal  (or  neutral)  sulphides,  such  as  NaHS  and 
Na2S. 

The  acid  sulphides  are  obtained  by  passing  the  gas  in  excess  into 
solutions  of  soluble  bases: 

H2S  +  NaOH  -»  H20  +  NaHS, 

and  are  neutral  in  reaction.  Their  negative  ion,  HS~,  gives  practically 
no  hydrogen-ion  (see  preceding  section).* 

By  adding  to  the  above  solution  an  amount  of  sodium  hydroxide 
equal  to  that  used  before,  and  driving  off  the  water  by  evaporation, 
the  second  unit  of  hydrogen  is  displaced,  and  normal  ("neutral") 
sodium  sulphide,  in  solid  form,  is  obtained: 

NaOH  +  NaHS  *=F  Na2S  +  H20  T  . 

This  action  is  wholly  reversed  when  dry  sodium  sulphide  is  dissolved 
in  water,  the  salt  being  completely  hydrolyzed  to  the  acid  salt: 

Na2S  ^  2Na+  -f-  S~~  /  < —  TJG— 

TT  /^  '       f~\~Cf—          TT4-      I    ^"*  -*"*-^    • 

*  In  point  of  fact,  N / 10  sodium-hydrogen  sulphide,  at  25°,  is  slightly  hydro- 
lyzed (0.14  per  cent,  James  Walker),  and  gives  therefore  a  faint  alkaline  reaction. 


SULPHUR  AND  HYDROGEN  SULPHIDE  419 

The  HS~  gives  a  lower  concentration  of  hydrogen-ion  than  does  the 
water,  and  hence  uses  up  in  its  formation  the  ions  of  hydrogen  pro- 
duced by  the  latter  until  an  amount  of  hydroxyl  equivalent  to  half 
the  sodium  is  formed.  The  abbreviated  equation  shows  this  more 
clearly: 

S=  +  H+  +  OH"  -» HS-  +  OH~. 

The  solution  is  therefore  strongly  alkaline  in  reaction.  In  general, 
a  normal  salt  derived  from  an  active  base  and  a  weak  acid  is  hydro- 
lyzed  to  some  extent  by  water  and  gives  an  alkaline  solution. 

In  the  abbreviated  formulation  used  above,  the  union  of  Na+ 
and  OH~  to  form  NaOH  is  not  shown,  because  it  is  slight  in  dilute 
solution  and  does  not  affect  the  result.  The  union  of  S=  and  H+ 
to  form  HS~  is  alone  shown,  because  it  is  extensive  and  significant. 
To  save  space,  this  plan  will  be  used  in  future,  where  the  same  situa- 
tion exists. 

The  soluble  acid  sulphides  are  oxidized  in  aqueous  solution  by 
atmospheric  oxygen: 

2NaSH  +  O2  ->  2NaOH  +  2S. 

The  sulphur  is  not  precipitated,  but  combines  with  the  excess  of  the 
sulphide,  forming  polysulphides  (see  below).  Some  sodium  thio- 
sulphate  is  produced  at  the  same  time. 

The  Action  of  Acids  on  Insoluble  Sulphides.  —  The  inter- 
action of  sulphides  and  acids  is  itself  so  important  a  matter  in  chemis- 
try, and  is  so  similar  in  theory  to  many  other  kinds  of  actions,  that 
special  attention  should  be  given  to  it.  The  common  method  of 
preparing  hydrogen  sulphide  from  ferrous  sulphide  affords  a  suitable 
illustration. 

Since  ferrous  sulphide  is  but  slightly  soluble  in  water,  the  action 
proceeds  by  a  rather  complex  series  of  equilibria: 

FeS  (solid)  ±?  FeS  (dslvd)  *=?  Fe++  +  S=  K 

2HC1  fc*  2Cr  +  2H+  |  ^  H2S  (dslvd)  ^  H2S  feas)- 

It  will  be  seen  that  a  number  of  reversible  changes  are  involved, 
and  the  question  is:  Why  does  the  reaction  proceed  forward,  as  it 
does?  To  answer  this  question,  a  consideration  of  each  of  the 
equilibria,  separately,  is  required. 

1.  The  dissolved  hydrogen  sulphide  is  very  feebly  ionized,  and 
maintains  a  smaller  concentration  of  sulphide-ion  S~  than  does 
ferrous  sulphide,  in  spite  of  the  comparative  insolubility  of  the 


420  INORGANIC  CHEMISTRY 

latter.  Hence,  the  S~  formed  from  the  FeS  is  continuously  re- 
moved by  union  with  the  hydrogen-ion  furnished  by  the  acid,  S=  + 
2H+  *=?  H2S,  and  all  the  other  equilibria  are  constantly  displaced 
forward  on  this  account.  The  action  is  therefore,  in  essence,  like 
neutralization  (p.  387). 

It  will  be  observed  that  the  action  takes  place  on  account  of  the 
feeble  ionization  of  the  weak  acid.  We  should  therefore  say  that  the 
weak  acid  withdraws,  and  not,  as  is  sometimes  done,  that  the  strong 
acid  drives  it  out. 

Since  the  action  is  an  ionic  one,  the  acids  must  be  employed  in 
dilute  form.  This  is  true  especially  of  oxygen  acids.  Thus,  concen- 
trated sulphuric  acid  has  little  action  upon  ferrous  sulphide  in  the 
cold,  and  when  the  substances  are  heated  the  oxygen  of  the  sulphuric 
acid  comes  into  play,  and  sulphur  dioxide  (q.v.)  and  free  sulphur  are 
formed. 

2.  The  union  of  S=  and  2H+  depends  on  the  magnitude  of  the 
product  of  their    concentrations   (p.   359),   [S=]  X  [H+]  X  [H+],    or 
[S=]  X  [H+]2.     Hence,  although  [S~]  is  minute,  on  account  of  the 
insolubility  of  FeS,  [H+]  is  large  on  account  of  the  great  dissociation 
of  the  HC1  and  the  fact  that  a  strong  solution  of  the  acid  can  be  used. 
Thus  the  product  may  be  large  enough  for  the  purpose. 

3.  When    a    still    less    soluble    sulphide,    like    cupric    sulphide 
CuS,  is  employed,  the  concentration  of  the  sulphide-ion  [S=]  is  too 
small  to  play  its  part  and  the  action  makes  almost  no  progress.     In 
this  case,  a  concentration  of  H+,  sufficient  to  raise  the  product  to 
the  necessary  value,  cannot  be  obtained  with  any  acid. 

4.  The  fact  that  hydrogen  sulphide  is  fairly  soluble  (3.6  vols.  :  1 
vol.)  hinders  the  action.     It  prevents  that  free  escape  of  one  prod- 
uct which  is  so  constantly  a  factor  in  promoting  reversible  chemical 
changes.     Thus,    if    cadmium    sulphide    CdS,    which   lies   between 
ferrous  and  cupric  sulphides  in  solubility,  is  employed  along  with 
rather  dilute  hydrochloric  acid,  a  concentration  of  hydrogen  sul- 
phide sufficient  to  stop  the  action  accumulates  before  the  liquid  is 
saturated  with  the  gas,  and  the  latter  can  begin  to  escape.     There 
are  then  two  ways  of  making  this  action  continuous.     Either  stronger 
hydrochloric  acid,  giving  a  higher  concentration  of  H+,  may  be  used 
to  force  the  formation  of  more  H2S  (by  union  of  2H+  and  S~),  or  the 
reverse  action,  due  to  accumulation  of  H2S  (dslvd),  may  be  diminished 
mechanically  by  leading  air  through  the  mixture  (p.  189)  and  so 
removing  the  hydrogen  sulphide  as  fast  as  it  is  formed.     Either  p)an 
will  cause  complete  interaction  with  the  cadmium  sulphide. 


SULPHUR  AND  HYDROGEN  SULPHIDE  421 

Classification  of  Insoluble  Sulphides.  —  In  analytical  chem- 
istry, advantage  is  taken  of  the  different  solubilities  of  the  sulphides, 
for  the  purpose  of  identifying  the  metallic  elements,  and  of  separating 
mixtures  containing  several  such  elements.  Three  classes  are  dis- 
tinguished. 

1.  The  sulphides  of  silver,   copper,  mercury,  and  some  other 
metals  are  exceedingly  insoluble,   and,  therefore,   do  not  interact 
with  dilute  acids  as  does  ferrous  sulphide  (p.  420).     These  may 
therefore  be  made  by  leading  hydrogen  sulphide  into  solutions  of 
their  salts: 

CuSO4  +  H2S  *=?  CuS  |  +  H2SO4. 

The  acid  produced  has  scarcely  any  effect  upon  the  sulphide,  and 
almost  no  reverse  action  is  observed.  In  this  action  the  sulphide- 
ion  is  the  active  substance  and,  by  its  removal,  all  the  equilibria 
are  displaced  forward. 

2.  The  sulphides  of  iron,  zinc,  and  certain  other  metals  are  insol- 
uble in  water,  but  not  so  much  so  as  the  last  class.     Hence  they  are 
decomposed  by  dilute  acids,  and  the  reverse  of  the  above  action 
takes  place  almost  completely.     These  sulphides  must  therefore  be 
made,  either  by  combination  of  the  elements,  or  by  adding  a  soluble 
sulphide  to  a  solution  of  a  salt: 

FeS04  +  (NH4)2S  fc*  FeS  j  +  (NH4)2S04. 

No  acid  is  produced  in  this  sort  of  interaction,  and  the  considerable 
insolubility  of  the  sulphide  of  iron  or  zinc  in  water  renders  the  change 
nearly  complete.  The  solubility  of  cadmium  sulphide  (p.  420) 
places  it  between  this  group  and  the  previous  one. 

3.  The  sulphides  of  barium,  calcium,  and  some  other  metals 
(q.v.),  although  insoluble  in  water,  are  hydrolyzed  by  it,  and  give 
soluble  products,  the  hydroxide  and  hydrosulphide: 

2CaS  +  2H20  <=*  Ca(OH)2  +  Ca(SH)2. 

They  may  be  prepared  by  direct  union  of  the  elements,  and  from 
the  sulphates  by  reduction  with  carbon.  But  they  are  not  pre- 
cipitated by  hydrogen  sulphide  or  ammonium  sulphide. 

Poly  sulphides.  —  When  sulphur  is  shaken  with  a  solution  of  a 
soluble  sulphide  or  acid  sulphide,  such  as  sodium  sulphide,  it  dis- 
solves, and  evaporation  of  the  solution  leaves  residues,  varying  in 
composition  from  Na^  to  Na2S5.  These  appear  to  be  mixtures 
composed  mainly  of  Na2S  and  Na2S4. 


422  INORGANIC  CHEMISTRY 

When  an  acid  is  poured  into  sodium  polysulphide  solution, 
minute  spherules  of  rhombic  sulphur  are  precipitated: 

Na2S4  +  2HC1  ->  2NaCl  +  H2S  f  +  3S  j  . 

This  precipitate  has  been  called  "soluble  amorphous  sulphur."  It  is 
certainly  all  soluble  in  carbon  disulphide,  but  the  particles  rotate  the 
plane  of  polarization  of  polarized  light  (A.  Smith),  and  are  therefore 
crystals.  Insoluble  amorphous  sulphur,  in  appreciable  quantities, 
cannot  be  obtained  under  any  conditions  from  polysulphides.  It  is 
obtained  by  adding  sodium  thiosulphate  solution  to  concentrated, 
active  acids. 

When  the  order  is  reversed,  sodium  pentasulphide  being  thrown 
into  concentrated  hydrochloric  acid,  no  hydrogen  sulphide  is  evolved. 
Hydrogen  pentasulphide,  H2S5,  a  yellow  oil,  falls  to  the  bottom  of  the 
vessel. 

The  Chemical  Relations  of  the  Element.  —  In  combination 
with  metals  and  hydrogen,  sulphur  is  bivalent,  forming  compounds 
like  H2S,  FeS,  CuS,  and  HgS.  In  combination  with  non-metals,  how- 
ever, the  valence  is  frequently  greater,  the  maximum  being  seen  in 
sulphur  trioxide,  where  we  must  assume  that  the  sulphur  is  sexivalent. 
Its  oxides  are  acid-forming,  and  it  is,  therefore,  a  non-metal. 

Sulphur  is  regarded  as  resembling  oxygen  more  closely  than  any 
of  the  other  elements  we  have  studied  so  far.  Both  unite  directly  with 
most  metals  and  non-metals.  In  this  they  are  like  chlorine.  But  hy- 
drogen chloride  is  highly  ionized  by  water,  while  the  hydrogen  com- 
pounds of  oxygen  and  sulphur  are  feebly  ionized.  The  formulae  of  the 
compounds  of  oxygen  and  sulphur  with  metals  are  similar,  CuO  and 
CuS,  NaOH  and  NaSH,  and  so  forth,  but  this  is  in  part  due  merely  to 
the  fact  that  both  elements  are  bivalent.  The  chemical  resemblance 
of  sulphur  to  selenium  and  tellurium  (q.v.)  is  much  more  striking  than 
its  resemblance  to  oxygen. 

Exercises.  —  1.  The  freezing-point  of  pure  sulphur  is  found  to 
vary  from  119°  down  to  114°,  depending  upon  the  temperature  to 
which  the  liquid  has  been  heated  and  the  speed  with  which  it  has  been 
cooled.  To  what  should  you  attribute  this  variability? 

2.  How  could  the  decomposition  of  hydrogen  sulphide  at  310°  be 
rendered  (a)  more  complete,  (6)  less  complete?  Would  the  percentage 
decomposed  be  affected  (a)  by  reducing  the  pressure,  (6)  by  mixing 
the  gas  with  an  indifferent  gas? 


SULPHUR  AND  HYDROGEN  SULPHIDE  423 

3.  What  are  the  relative  volumes  of  the  gases  (p.  259)  in  the 
action  of,  (a)  hydrogen  iodide  and  sulphur,  (6)  hydrogen  sulphide 
and  sulphur  dioxide? 

4.  To  what  classes  of  ionic  actions  (p.  402)  do  the  interactions  of 
hydrogen  sulphide  solution  and  (a)  oxygen  (p.  417),  (b)  acidified 
potassium  dichromate  (p.  418),  (c)  sodium  hydroxide  (p.  418),  (d) 
iodine  (p.  278),  belong? 

5.  Why  is  normal  sodium  sulphide  only  half  hydrolyzed  by  water? 

6.  Formulate  completely,  after  the  model  on  p.  419,  the  actions 
of,  (a)  hydrogen  sulphide  and  cupric  sulphate  solution;    (6)   am- 
monium sulphide  and  ferrous  sulphate.     In  each  case  explain  which 
equilibrium  determines  the  direction  of  the  action. 


CHAPTER  XXI 
THE   OXIDES  AND   OXYGEN  ACIDS   OF   SULPHUR 

FouR  oxides  of  sulphur,  represented  by  the  formulae  82(1)3,  SO2, 
80s,  and  8207,  are  known.  Of  these,  however,  the  first  and  the  last 
are  much  less  familiar  substances  than  the  other  two.  The  dioxide 
and  trioxide  of  sulphur  are  not  only  important  in  themselves,  but 
their  relation  to  the  acids  H2SO3  and  H2SO4,  which  may  be  obtained 
from  them  by  the  addition  of  water,  makes  them  doubly  so  to  the 
chemist. 

The  Preparation  of  Sulphur  Dioxide  SO2.  —  1.  When  sul- 
phur burns  in  air  or  oxygen,  sulphur  dioxide  is  produced  (p.  88). 

2.  The  larger  part  of  the  sulphur  dioxide  used  in  commerce  is 
probably  obtained  by  the  roasting  (calcining)  of  sulphur  ores.     Pyrite 
FeS2,  for  example,  which  is  a  familiar  yellow,  metallic-looking  mineral, 
can  be  burnt  in  a  suitable  furnace  on  account  of  the  large  amount  of 
sulphur  which  it  contains: 

4FeS2  +  HOj  ->  2Fe203  +  8S02  T  . 

The  gas,  although  mixed  with  a  great  amount  of  nitrogen  which 
entered  as  part  of  the  air,  can  be  used  to  make  sulphuric  acid. 

It  should  be  noted,  in  passing,  that  heating  and  roasting  or  cal- 
cining are  distinct  processes  in  chemistry.  Roasting  or  calcining 
always  assumes  the  access  of  the  air  and  employment  of  its  oxygen; 
heating,  in  the  absence  of  modifying  words,  assumes  the  exclusion  or 
the  chemical  indifference  of  the  air. 

3.  In  the  laboratory,  a  steady  stream  of  gas  is  obtained  by 
allowing  hydrochloric  acid  to  drop   upon  solid   sodium  acid  sul- 
phite, or  concentrated  sulphuric  acid  to  trickle  into  a  40  per  cent 
solution  of  the  same  salt  (Fig.  41)  : 

°3 


The  sulphurous  acid  being  only  moderately  ionized,  its  molecules  are 
formed  in  considerable  amount.     Being  also  unstable,  it  decomposes 

424 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         425 

spontaneously  into  water  and  sulphur  dioxide,  and  the  latter  escapes 
when  sufficient  water  for  its  solution  is  not  present. 

4.  Sulphur  dioxide  can  also  be  made  by  the  reduction  of  con- 
centrated sulphuric  acid  by  copper  at  a  high  temperature.  Copper  is 
the  metal  most  commonly  employed,  because  by  its  means  very  pure 
sulphur  dioxide  can  be  obtained.  More  active  metals,  such  as  iron 
and  zinc,  although  cheaper,  cannot  be  used,  since  they  reduce  the 
sulphuric  acid  to  hydrogen  sulphide.  The  undiluted  hydrogen  sul- 
phate consists  entirely  of  molecules  and,  at  the  high  temperatures  at 
which  alone  the  action  is  vigorous,  is  an  oxidizing  agent  (cf.  p.  119). 
A  part  of  the  acid  loses  oxygen  to  form  water  with  the  hydrogen 
of  another  molecule: 

H2S04  ->  H20  +  S02  (+  O)  (1) 

(O)  +  H2SO4  +  Cu  -^  H2O  +  CuSO4  (2) 

2H2SO4  +  Cu  ->  2H2O  +  SO2  +  CuSO4 

Some  easily  oxidized  non-metals,  such  as  carbon  and  sulphur,  act  in 
the  same  way: 

C  +  2H2S04  -+  2H2O  +  2S02  -f-  CO2, 
S  +  2H2SO4  -»  2H2O  +  3SO2. 

Making  Equations  by  Positive  and  Negative  Valences.  — 

Equations  like  the  foregoing  can  be  constructed  also  by  assuming 
that  each  element  in  a  compound  is  either  positive  or  negative, 
and  by  marking  the  valences  accordingly  (for  details,  see  p.  493). 
Thus,  in  sulphuric  acid,  we  have  2H+  (positive,  univalent)  and  4O= 
(each  bivalent  and  negative).  Since  the  numbers  of  positive  and 
negative  valences  must  be  equal,  and  we  have  2®*  and  80,  it 
follows  that  the  sulphur  carries  6©,  S+++. 

Now  when,  in  making  the  experiment,  we  find  the  products  SO2 
and  CuSO4,  we  may  infer  that  the  hydrogen  formed  water.  We 
infer,  also,  that  to  obtain  two  compounds  containing  sulphur,  at 
least  2H2SO4  was  required.  We  then  note  that  the  S  in  S02  is  quad- 
rivalent. Hence  S+++  became  S++  and  2©  were  released.  The 
metallic  copper  used  was  free  and  without  valence,  and  became 
CuS04,  in  which  it  is  Cu4"*".  It  obtained  the  2©  from  the  sulphur. 
The  action  can  therefore  be  analyzed  as  follows: 

*  The  signs  ©  and  ©  stand  for  quantities  of  electricity  equal  to  those  car- 
ried by  one  equivalent  of  an  ionic  substance,  and  therefore  required  for  its  dis- 
charge and  liberation. 


426  INORGANIC  CHEMISTRY 

[2H+  +  StS  +  40=]  -»  [Stt  +  20=]  +  [2H+  +  O=]  +  O=  -f  20 
First  H2S04  S02  H20  Balance 

The  second  H2SO4  gives  [2H+  +  SO4=].  The  Cu  takes  the  20  giving 
Cu++,  and  this  with  the  S04=  gives  CuS04.  The  2H+  takes  the  O= 
from  the  balance,  giving  H2O.  Thus,  the  whole  balance  is  used  and 
the  products  are  accounted  for.  The  equation  must  therefore  be: 

2H2S04  +  Cu  -*  S02  +  2H2O  +  CuS04. 

It  will  be  noted  that  the  two  molecules  of  sulphuric  acid  play  different 
roles.  Only  one  of  them  is  used  in  oxidizing. 

Similarly,  with  sulphuric  acid  and  carbon,  the  same  analyzed 
equation  applies.  The  carbon  gives  C02.  Thus,  the  carbon  goes 
from  C°  to  C++.  To  obtain  the  40,  2H2SO4  is  required  (equation 
above).  Hence, 

2H2S04  +  C  ->  C02  +  2SO2  +  2H2O. 

When  hydrogen  sulphide  is  led  through  concentrated  sulphuric 
acid,  the  latter  is  reduced  to  sulphur  dioxide,  and  the  former  is 
oxidized,  giving  free  sulphur  (p.  416) : 

2H+  +  S=  +  2©  ^2H+  +  S°  |  . 

Since  this  action  requires  20,  and  sulpnuric  acid  in  giving  SO2 
delivers  20,  it  follows  that  1H2S04  will  decompose  1H2S: 

H2S04  +  H2S  -»  2H20  +  S°  |  +  S02. 

Finally,  when  HI  with  sulphuric  acid  (p.  277)  gives  free  iodine 
(1°),  and  H2S  (2H+  +  S=),  evidently  SK+  in  sulphuric  acid  gives  up 
80,  becoming  S~: 

[2H+  +  STS  +  40=]  -»  2H+  +  S=  +  40=  +  8  © 
and  [H+  +  I"]  +  ©  ->  H+  +  1°. 

Evidently,  1H2S04  giving  8©  will  interact  with  SHI,  changing  81" 
into  81°.  Hence, 

H2S04  +  SHI  ->  4H2O  +  H2S  +  81°. 

The  reader  should  practice  the  use  of  this  method  by  making  the 
equations  for  the  actions  of  zinc  (p.  119)  and  of  hydrogen  bromide 
(p.  272)  upon  sulphuric  acid. 

Physical  Properties.  —  Sulphur  dioxide  is  a  gas  possessing  a 
penetrating  and  characteristic  odor.  This  is  frequently  spoken  of  as 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR 


427 


FIG.  111. 


the  "odor  of  sulphur,"  but  it  should  be  remembered  that  sulphur 
itself  has  scarcely  any  smell  at  all.  The  weight  of  the  G.M.V.  of 
the  gas  (65.54  g.)  shows  it  to  be  more  than  twice  as  heavy  as  air.  The 
critical  temperature  is  156°.  By  means  of  a  _^, 
freezing  mixture  of  ice  and  salt  (Fig.  Ill),  the 
gas  is  easily  condensed  in  a  U-tube  to  a  trans- 
parent, mobile  fluid,  which  boils  at  —8°.  At 
20°  the  vapor  tension  of  the  liquid  is  3.25  at- 
mospheres, so  that  the  liquid  is  handled  and 
sold  in  glass  syphons,  or  in  sealed  tin  cans.  The 
liquid  may  be  frozen  to  a  white  solid,  melting  at 
—  76°.  It  ionizes  substances  dissolved  in  it  as 
well  as  does  water  (cf.  p.  357).  The  solubility 
of  the  gas  in  water,  5000  volumes  in  100,  is  very 
great.  Unlike  solutions  of  the  hydrogen  halides 
(p.  211),  however,  the  liquid  is  completely  freed  from  the  gas  by 
boiling. 

.      *m 

Chemical  Properties.  —  Sulphur  dioxide  is  stable,  being  de- 
composed only  by  the  use  of  a  very  high  temperature. 

It  unites  with  water  to  form  sulphurous  acid,  H2S03,  which  is 
unstable  and  exists,  only  in  solution.  Although  the  gas  itself  some- 
times receives  this  name,  it  is  not  acid:  it  is  simply  the  anhydride 
(p.  150)  of  the  acid.  The  only  compound  of  SO2  and  H2O  that  has 
been  isolated  is  a  solid  hydrate  SO2,7H2O. 

Since  the  maximum  valence  of  sulphur  is  6,  sulphur  dioxide,  in 
which  but  four  of  the  valences  of  sulphur  are  used,  is  unsaturated. 
It  is  therefore  still  able  to  combine  directly  with  suitable  elements, 
such  as  chlorine  and  oxygen.  When  it  is  mixed  with  chlorine  in  sun- 
light, a  liquid,  sulphuryl  chloride  SO2C12,  is  produced. 

Liquefied  sulphur  dioxide  is  now  sold  in  tin  cans,  and  is  employed 
for  bleaching  straw,  wool,  and  silk.  As  a  disinfectant  it  has  been 
displaced  to  a  large  extent  by  formaldehyde. 

The  Liquefiabilities  of  Gases,  —r  It  will  assist  us  in  recalling 
which  gases  are  hard  to  liquefy  and  which  easy,  if  we  memorize 
the  fact  that  Faraday  (from  1823  to  1845)  liquefied  most  of  the  fa- 
miliar gases  and  failed  only  with  three,  namely  hydrogen  (c.t.  —  242°), 
oxygen  (c.t.  -113°),  and  nitrogen  (c.t.  -146°).  These,  with  nitric 
oxide  NO  (c.t.  -93.5°),  carbon  monoxide  CO  (ct.-40°),  methane 
CH4  (c.t.  —99°),  and  the  six  inert  gases  (p.  9),  are  the  ones  which 


428  INORGANIC  CHEMISTRY 

have   low  critical  temperatures    (cf.  p.   166)    and  are  difficult  to 
liquefy. 

Of  the  gases  we  have  studied,  the  ones  which  are  more  or  less 
easily  liquefied  are :  hydrogen  chloride  (c.t.  +52°),  bromide,  and  iodide, 
chlorine  (c.t.  +141°),  ozone,  hydrogen  sulphide  (c.t.  +100°),  sulphur 
dioxide  (c.t.  +154°). 

The  Solubilities  of  Gases.  —  For  the  purpose  of  remembering 
the  solubilities  of  gases  in  water,  it  is  convenient  to  divide  the  gases 
into  three  classes.  The  following  are  the  ones  we  have  studied: 

1.  Slightly  soluble:  Oxygen  (4  vol.  :  100  at  0°),  hydrogen  (2  :  100 
at  0°). 

2.  Soluble:    Chlorine   (2.6  vol.  :  1   at   10°),    hydrogen    sulphide 
(4.4  : 1  at  0°). 

3.  Very  soluble:  Hydrogen  chloride   (505  vol.  :  1  at  0°),  bromide 
(404  :  1)  and  iodide  (1570  :  1),  sulphur  dioxide  (69  :  1  at  0°). 

Preparation  of  Sulphur  Trioxide  SO3.  —  Although  the  forma- 
tion of  sulphur  trioxide  S0§  is  accompanied  by  the  liberation  of  much 
heat,  sulphur  dioxide  and  oxygen,  even  when  heated  together,  unite 
very  slowly.  Ozone,  however,  combines  with  the  former  readily. 

The  interaction  of  sulphur  dioxide  and  oxygen  is  hastened  by 
many  substances,  such  as  glass,  porcelain,  ferric  oxide  and,  more 
especially,  finely  divided  platinum,  which  remain  themselves  un- 
changed and  simply  act  as  catalytic  agents.  The  contact  process, 
as  this  is  called,  has  been  rendered  available  for  the  commercial 
manufacture  of  sulphur  trioxide  by  Knietsch  (1901).  The  chief 
features  of  the  process  are:  (1)  The  complete  removal  of  arsenious 
oxide,  dust,  and  other  impurities  derived  from  the  calcining  of  pyrite 
or  some  other  mineral  sulphide,  the  minutest  traces  of  which  "  poison  " 
the  catalytic  agent  and  soon  render  it  absolutely  inoperative.  (2) 
The  preliminary  passage  of  the  cold  mixture  of  gases  over  the  outside 
of  the  pipes  containing  the  contact  agent.  This  removes  part  of  the 
heat  generated  by  the  action,  SO2  +  O  <=±  SO3  +  22,600  cal.,  going 
on  inside,  and  keeps  the  temperature  of  the  interior  at  400°.  Below 
400°,  the  union  is  too  slow;  above  400°,  the  reverse  action  is  strength- 
ened (van't  Hoff's  law,  p.  305),  and  the  union  is  incomplete.  At  400°, 
98-99  per  cent  of  the  materials  unite;  at  700°,  only  60  per  cent,  at 
900°  practically  none.  Twice  the  quantity  of  oxygen  theoretically 
needed  is  employed. 

The  vaporous  product,  mainly  1  vol.  02  to  2  vol.  SO3  (gas),  is 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         429 

condensed  by  being  led  into  97-99  per  cent  sulphuric  acid,  and  the 
concentration  of  the  liquid  is  constantly  maintained  at  this  point  by 
the  regulated  influx  of  water.  Oleum  (q.v.)  is  also  made  by  omitting 
to  add  water.  The  trioxide  is  thus  chiefly  used  for  immediate  con- 
version into  sulphuric  acid. 

The  process  may  be  illustrated  by  placing  some  platinized  as- 
bestos* in  a  tube  (Fig.  88,  p.  217),  which  is  gently  warmed,  and  in- 
troducing oxygen  and  sulphur  dioxide  through  the  limbs  of  the 
Y-tube.  Dense  fumes  appear  at  the  exit  (see  next  par.). 

It  would  seem  to  be  simpler  to  dissolve  the  gaseous  sulphur  tri- 
oxide in  water,  to  give  sulphuric  acid  H2O  +  SO3  —  >  H2SO4,  rather 
than  in  98  per  cent  sulphuric  acid,  but  this  cannot  be  done.  The 
mixture  O2  +  2SO3  is  very  incompletely  absorbed  by  water.  When 
a  bubble  of  this  mixture  enters  water,  the  latter  evaporates  in  the 
attempt  to  saturate  the  space  occupied  by  the  bubble  with  water 
vapor  (p.  145  and  Appendix  IV).  The  water,  however,  combines 
with  the  sulphur  trioxide  to  form  a  fog,  consisting  of  droplets  of  liquid 
sulphuric  acid,  and  so  more  and  more  water  evaporates  into  the 
bubbles.  The  molecules  of  80s,  so  long  as  they  remain  gaseous,  move 
with  great  velocity,  namely  292  meters  per  second  at  room  tempera- 
ture, and  still  faster  in  this  hot  gaseous  mixture.  Hence,  all  the 
molecules  that  escape  combination  with  the  water  vapor,  strike  the 
wall  of  the  bubble,  and  combine  with  the  water  in  a  few  seconds.  The 
droplets  of  sulphuric  acid,  forming  the  fog,  however,  are  not  mole- 
cules but  large  aggregates  of  molecules.  They  do  not  therefore  move 
like  the  molecules  of  a  gas,  but  are  absolutely  stationary.  Hence, 
after  the  gaseous  sulphur  trioxide  has  dissolved,  the  droplets  of  fog, 
carried  by  the  excess  of  oxygen,  can  be  bubbled  through  a  whole  series 
of  vessels  of  water  in  succession  without  any  appreciable  number  of 
the  droplets  being  dissolved.  The  same  fog  can  be  shaken  in  a  flask 
with  water,  violently  and  continuously,  without  any  appreciable 
solution.  When  the  water  is  thrown,  by  the  shaking,  through  the 
oxygen,  the  oxygen  is  split  up  by  the  water,  and  driven  about,  but  the 
fog  particles  move  with  the  oxygen,  so  that  the  water  never  reaches 
them.  On  the  other  hand,  when  the  mixture  of  gases  bubbles 
through  98  per  cent  sulphuric  acid,  as  is  done  in  practice,  there  is  no 
water  available,  the  sulphur  trioxide  remains  gaseous,  and  its  rapidly 
moving  molecules  in  a  few  seconds  have  all  plunged  into  the  sulphuric 
acid  and  combined  with  it,  either  uniting  with  the  1-3  per  cent  of 


*  Asbestos,  dipped  in  a  solution  of  chloroplatinic  aci4  ftftd  heated  in  the 
Bunsen  flame:  HzPtCU  --»  Pt  +  2HC1  T  +  2C12  T. 


430  INORGANIC  CHEMISTRY 

water  present,  or,  when  oleum  is  made,  uniting  with  the  sulphuric 
acid  to  form  pyrosulphuric  acid:  S03  -f  H2S04  — >  H2S2O7. 

This  case  affords  an  admirable  illustration  of  the  importance  of 
physics  in  practical  chemistry  (p.  40).  The  chemical  reaction  occurs 
with  water,  but  the  physical  condition  of  the  fog  of  sulphuric  acid 
prevents  its  dissolving  and,  if  water  were  used  in  a  factory,  a  large 
proportion  of  the  sulphuric  acid  would  pass  with  the  excess  of  oxygen 
into  the  air  and  be  lost.  In  fact,  it  would  kill  vegetation,  and  make 
life  unbearable  in  the  neighborhood.  It  is  stated  that  the  inventor 
of  the  contact  process  spent  a  year  of  time,  and  much  money,  in 
trying  to  find  some  way  of  using  water  to  absorb  the  gas.  Yet  ten 
minutes'  consideration  of  the  physical  situation  would  have  shown 
that  this  was  impossible,  and  the  rest  of  the  year  could  have  been 
devoted  to  other  work  that  offered  some  prospect  of  successful 
results. 

Formerly  sulphur  trioxide  was  obtained  by  the  distillation  of 
impure  ferric  sulphate,  Fe2(SO4)3  — >  Fe2O3  +  3SO3.  It  may  also  be 
prepared  by  repeated  distillation  of  concentrated  sulphuric  acid  with 
a  powerful  drying  agent,  like  phosphoric  anhydride. 

Physical  Properties.  —  Sulphur  trioxide  S03  is,  at  ordinary 
temperatures,  fluid.  The  crystals,  obtained  by  cooling,  melt  at  14.8°. 
The  liquid  boils  at  46°,  and  is,  therefore,  exceedingly  volatile  at  ordi- 
nary temperature.  It  fumes  strongly  when  exposed  to  the  air,  in  con- 
sequence of  the  union  of  the  vapor  with  moisture  and  the  production 
of  minute  drops  of  sulphuric  acid. 

A  white  crystalline  variety  of  the  substance,  which  in  appearance 
closely  resembles  asbestos,  is  obtained  when  a  trace  of  water  has 
gained  access  to  the  oxide.  The  substance  is  dimorphous  (p.  412). 
When  heated  to  50°,  this  passes  into  vapor  of  SO3  without  melting. 
This  white  solid  is  the  more  stable  and  more  familiar  form  of  the 
trioxide. 

Chemical  Properties.  —  The  vapor  of  sulphur  trioxide,  when 
heated,  dissociates  into  sulphur  dioxide  and  oxygen  (400°,  2  per  cent; 
700°,  40  per  cent). 

Sulphur  trioxide  is  not  itself  an  acid,  but  it  is  the  anhydride  of 
sulphuric  acid.  When  placed  in  water  it  unites  vigorously,  causing 
a  hissing  noise  due  to  the  steam  produced  by  the  heat  of  the  union. 
In  consequence  of  its  great  tendency  to  combine  with  water,  the  liquid 
variety,  which  is  the  more  active,  removes  the  elements  of  this  sub- 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         431 

stance  from  materials  which  contain  them  in  the  proper  proportions. 
Thus  paper,  which  is  largely  cellulose  (CeHioOs)^,  and  sugar  C^H^On, 
are  charred  by  it,  and  carbon  is  set  free. 

Just  as  sulphur  trioxide  unites  with  water  to  give  hydrogen  sul- 
phate, so  it  combines  vigorously  with  many  oxides  of  metals,  produ- 
cing the  corresponding  sulphates: 

H20  +  S03  *=>  H2S04,        CaO  +  S03  -»  CaS04. 

The  union  of  an  oxide  of  a  non-metal  with  the  oxide  of  a  metal,  in 
this  fashion,  is  a  general  method  of  obtaining  salts  (cf.  p.  407). 

Oxygen  Acids  of  Sulphur.  —  Sulphurous  and  sulphuric  acids 
have  been  mentioned  frequently  already.  Next  to  them  in  importance 
come  thiosulphuric  acid  and  persulphuric  acid.  The  compositions  of 
the  acids  show  their  relationships  : 


Hyposulphuro^s  acid,  H2S204.  Sodium  hyposulphite, 

Sulphurous  acid,  H2SO3.  Sodium  sulphite,  Na2SOa. 

Sulphuric  acid,  H2S04.  Sodium  sulphate,  Na2S04. 

Thiosulphuric  acid,  H2S203.  Sodium  thiosulphate,  Na2S203. 

Persulphuric  acid,  H2S2Os.  Sodium  persulphate,  Na2S2Os. 

Thiosulphuric  acid  (Gk.  Ouov,  sulphur)  is  so  named  because  it  con- 
tains one  unit  of  sulphur  in  place  of  one  of  the  units  of  oxygen  of 
sulphuric  acid.  Note  that  when  the  names  of  the  acids  end  in  ous 
and  ic,  the  names  of  the  salts  end  in  He  and  ate,  respectively.  Besides 
the  above  we  have  also  the  poly  thionic  acids,  namely  :  Dithionic  acid 
H2S2Oe,  trithionic  acid  H2S3Oe,  tetrathionic  acid  H^Oe,  and  penta- 
thionic  acid  H^SsOe. 

On  account  of  its  commercial  importance  and  the  interest  attach- 
ing to  its  method  of  manufacture  and  to  its  properties,  we  may  first 
discuss  sulphuric  acid.  We  shall  then  be  able  to  dispose  of  the  re- 
maining acids  in  a  much  briefer  fashion. 

SULPHURIC  ACID  H2S04 

Although  salts  of  sulphuric  acid,  such  as  calcium  sulphate,  are  ex- 
ceedingly plentiful  in  nature,  the  preparation  of  the  acid  by  chemi- 
cal action  upon  the  salts  is  not  practicable.  The  sulphates,  indeed, 
interact  with  all  acids,  but  the  actions  are  reversible.  The  comple- 
tion of  the  action  by  the  plan  used  in  making  hydrogen  chloride 
(p.  206),  involving  the  removal  of  the  sulphuric  acid  by  distillation, 


432  INORGANIC  CHEMISTRY 

would  be  difficult  on  account  of  the  involatility  of  this  acid.  It  boils 
at  330°;  and  suitable  active  acids,  less  volatile  still,  which  might  be 
used  to  liberate  it,  do  not  exist.  We  are  therefore  compelled  to  build 
up  sulphuric  acid  from  its  elements. 

The  union  of  sulphur  dioxide  and  oxygen  by  the  contact  process, 
and  combination  of  the  trioxide  with  water  (p.  428),  is  the  best  method 
for  making  a  highly  concentrated  acid.  For  obtaining  ordinary  "oil 
of  vitriol/'  however,  the  chamber  process  is  still  used  extensively. 

History  of  Sulphuric  Acid  Manufacture.  —  Impure  forms  of 
sulphuric  acid  have  been  known  for  many  centuries.  In  the  fifteenth 
century  it  was  made  by  distilling  ferrous  sulphate  with  sand.  The 
product,  however,  contained  much  water  and  sulphur  dioxide.  The 
first  successful  preparation  of  the  substance  commercially  was  made 
by  Ward  at  Richmond-on-the-Thames  (1758) .  The  process  consisted 
in  burning  a  mixture  of  sulphur  and  saltpeter  KNOs  in  a  ladle 
suspended  in  a  large  glass  globe  partially  filled  with  water.  The 
gases  which  were  evolved  contained  large  quantities  of  sulphur 
dioxide  and  oxides  of  nitrogen,  which,  by  interaction  with  atmospheric 
oxygen  and  water  (see  below),  produced  the  sulphuric  acid.  The  solu- 
tion which  was  obtained,  although  it  could  be  prepared  of  any  desired 
concentration  by  the  burning  of  a  sufficient  number  of  charges,  was 
far  from  pure  and  was  expensive,  bringing  thirteen  shillings  ($3.25) 
per  pound.  Subsequently  a  chamber  lined  with  lead  was  substituted 
for  the  glass  vessel.  This  reduced  the  price  to  about  two  shillings 
and  sixpence  ($0.60)  per  pound.  The  same  principles  are  used  in 
the  modern  "  chamber  process." 

Chemistry  of  the  Chamber  Process.  —  The  gases,  the  inter- 
actions of  which  result  in  the  formation  of  sulphuric  acid,  are :  water 
vapor,  sulphur  dioxide,  nitrous  anhydride  N2Oa*  (q.v.),  and  oxygen. 
These  are  obtained,  the  first  by  injection  of  steam,  the  second  usually 
by  the  burning  of  pyrite,  pyrrotite  (FeS),  or  some  other  mineral 
sulphide,  the  third  from  nitric  acid  HNOa,  and  the  fourth  by  the 
introduction  of  air.  The  gases  are  thoroughly  mixed  in  large  leaden 
chambers,  and  the  sulphuric  acid  condenses  and  collects  upon  the 
floors.  In  spite  of  elaborate  investigations,  instigated  by  the  exten- 

*  This  gas  is  unstable,  breaking  up  in  part  into  nitric  oxide  NO  and  nitrogen 
tetroxide  NO2:  N2O3  <=*  NO  +  NO2.  In  the  process  here  discussed,  however, 
the  mixture  behaves  as  if  it  were  all  NzOz,  and  so  only  nitrous  anhydride  is  named 
in  this  connection,  , 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         433 

sive  scale  upon  which  the  manufacture  is  carried  on  and  the  immense 
financial  interests  involved,  some  uncertainty  still  exists  in  regard  to 
the  precise  nature  of  the  chemical  changes  which  take  place.  Accord- 
ing to  Lunge,  supporting  the  view  first  suggested  by  Berzelius,  the 
greater  part  of  the  product  is  formed  by  two  successive  actions,  the 
first  of  which  yields  a  complex  compound  that  is  decomposed  by 
excess  of  water  in  the  second: 

/0-H 

H2O  +  2S02  +  N2O3  +  02  ->  2SO2  '  (1) 

NO  —  NO 

The  group  —NO  is  found  in  many  compounds.  Here,  if  it  were  dis- 
placed by  hydrogen,  sulphuric  acid  would  result.  Hence  this  com- 
pound is  called  nitrosylsulphuric  acid. 

/0-H  _^          /OH 

\  n Ain  \  QJJ 


The  equations  (1)  and  (2)  are  not  partial  equations  for  one  inter- 
action, but  represent  distinct  actions  which  can  be  carried  out  sepa- 
rately. In  a  properly  operating  plant,  indeed,  the  nitrosylsulphuric 
acid  is  not  observed.  But  when  the  supply  of  water  is  deficient,  white 
"  chamber  crystals/'  consisting  of  this  substance,  collect  on  the  walls. 

The  explanation  of  the  success  of  this  seemingly  roundabout 
method  of  getting  sulphuric  acid  is  as  follows:  The  direct  union  of 
sulphur  dioxide  and  water  to  form  sulphurous  acid  is  rapid,  but  the 
action  of  free  oxygen  upon  the  latter,  2H2S03  +  O2  — >  2H2SO4,  is  ex- 
ceedingly slow.  Reaching  sulphuric  acid  by  the  use  of  these  two 
changes,  although  they  constitute  a  direct  route  to  the  result,  is  not 
feasible  in  practice.  On  the  other  hand,  both  of  the  above  actions, 
(1)  and  (2),  happen  to  be  much  more  speedy,  and  so,  by  their  use, 
more  rapid  production  of  the  desired  substance  is  secured  at  the  ex- 
pense of  a  slight  complexity.  It  may  be  added  that  the  heat  finally 
given  out  in  the  formation  of  one  formula-weight  of  sulphuric  acid  is 
exactly  the  same  in  amount  whether  nitrous  anhydride  intervenes  or 
not  (c/.  p.  100). 

The  progress  of  the  first  action  is  marked  by  the  disappearance  of 
the  brown  nitrous  anhydride  and,  on  the  introduction  of  water,  the 
completion  of  the  second  results  in  the  reproduction  of  the  same  sub- 
stance. It  would  thus  seem  as  if  the  nitrous  anhydride  should  take 
part  an  indefinite  number  of  times  in  these  changes  and  so  facilitate 
the  conversion  of  an  unlimited  amount  of  sulphur  dioxide,  oxygen, 


434  INORGANIC   CHEMISTRY 

and  water  into  sulphuric  acid,  without  diminution  of  its  quantity. 
In  practice,  however,  certain  subsidiary  actions  take  place,  such  as, 
for  example,  the  reduction  of  some  nitrous  anhydride  to  nitrous  oxide 
N2O,  which  permanently  remove,  a  part  of  the  material  from  par- 
ticipation in  the  cycle, 

The  loss  of  nitrous  anhydride  is  made  good  by  the  introduction 
of  nitric  acid  vapor  into  the  chamber.  This  acid  is  secured  by  the 
action  of  concentrated  sulphuric  acid  upon  commercial  sodium  nitrate 
NaN03: 

NaN03  +  H2S04  *±  HNO3 1  +  NaHSO4. 

On  account  of  the  volatility  of  the  nitric  acid,  a  moderate  heat  is 
sufficient  to  remove  it  from  admixture  with  the  other  substances,  and 
its  vapor  is  swept  along  with  the  other  gases  into  the  apparatus.  The 
initial  action  which  the  nitric  acid  undergoes: 

H20  +  2SO2  +  2HNO3  -*  2H2SO4  +  N203 
may  be  written  to  show  the  anhydride  of  nitric  acid:     . 
H20  +  2S02  +  H20,N205  -» 2H2S04  +  N2O3. 

The  two  molecules  of  water,  one  actually,  the  other  potentially,  pres- 
ent, with  the  two  molecules  of  sulphur  dioxide,  can  furnish  two 
molecules  of  sulphurous  acid  H2S03.  The  N2Os  in  passing  to  the 
condition  N2O3  gives  up  the  two  units  of  oxygen  required  to  convert 
this  sulphurous  acid  into  sulphuric  acid. 

Details  of  the  Chamber  Process.  —  The  sulphur  dioxide  is 
produced  in  a  row  of  small  furnaces  A  (Fig.  112),  the  structure  of 
which  depends  upon  the  nature  of  the  substance  employed  to  yield 
this  fundamental  constituent  of  sulphuric  acid.  When  good  pyrite 
is  used,  the  ore  burns  unassisted  (p.  424),  while  impure  pyrite  and 
zinc-blende  ZnS  have  to  be  heated,  to  a  greater  or  less  degree,  arti- 
ficially, to  maintain  the  combustion.  The  gases  from  the  various 
furnaces  pass  into  one  long  dust  flue,  in  which  they  are  mingled  with 
the  proper  proportion  of  air,  and  have  an  opportunity  to  deposit 
oxides  of  iron  and  of  arsenic  and  other  materials  which  they  transport 
mechanically.  From  this  flue  they  enter  the  Glover  tower  G,  in  which 
they  acquire  the  oxides  of  nitrogen.  Having  secured  all  the  necessary 
constituents,  excepting  water,  and  having  been  reduced  very  con- 
siderably in  temperature,  the  gases  next  enter  the  first  of  the  lead 
chambers,  large  structures,  from  three  to  five  in  number,  lined  com- 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR 


435 


pletely  with  sheet  lead.  They  vary  in  size,  measuring  as  much  as  100 
X  40  X  40  feet,  and  sometimes  having  a  total  capacity  of  150,000  to 
200,000  cubic  feet.  As  the  gases  drift  through  these  chambers  they 
are  thoroughly  mixed,  and  an  amount  of  water  considerably  in  excess 
of  that  required  by  the  chemical  reactions,  is  injected  in  the  form  of 
steam  at  various  points.  The  temperature  in  the  first  chamber  is 


FIG.  112. 


maintained  at  50°  to  65°,  while  in  the  last  chamber  it  is  about  15° 
above  that  of  the  outside  air.  The  acid,  along  with  the  excess  of 
water,  condenses  and  collects  upon  the  floor  of  the  chamber,  while  the 
unused  gases,  consisting  chiefly  of  nitrous  anhydride  and  a  very  large 
amount  of  nitrogen,  derived  from  the  air  originally  admitted,  find  an 
exit  into  the  Gay-Lussac  tower  L. 

This  is  a  tower  about  fifty  feet  in  height,  filled  with  tiles,  over 
which  concentrated  sulphuric  acid  continually  trickles  from  a  reservoir 
at  the  top.  The  object  of  this  tower  is  to  catch  the  nitrous  anhydride 
and  enable  it  to  be  reemployed  in  the  process.  This  is  accomplished 
by  a  reversal  of  action  (2)  above.  The  acid  which  accumulates  in  the 
vessel  at  the  bottom  of  this  tower  contains  the  nitrosylsulphuric  acid, 
and  by  means  of  compressed  air  is  forced  through  a  pipe  up  to  a 
vessel  at  the  top  of  the  Glover  tower  G.  When  this  " nitrous  vitriol" 
is  mixed  with  dilute  sulphuric  acid  from  a  neighboring  vessel,  by 
allowing  both  to  flow  down  into  the  tower,  the  nitrous  anhydride  is 
once  more  set  free  by  the  interaction  of  the  water  in  the  dilute  acid 


436  INORGANIC  CHEMISTRY 

(action  (2)).  The  Glover  tower  is  filled  with  broken  quartz  or  tiles, 
and  the  heated  gases  from  the  furnace  acquire  in  it  their  supply  of 
nitrous  anhydride.  Their  high  temperature  causes  a  considerable 
concentration  of  the  diluted  sulphuric  acid  as  it  trickles  downward. 
The  acid,  after  traversing  this  tower,  is  sufficiently  strong  to  be  used 
once  more  for  the  absorption  of  nitrous  anhydride. 

To  replace  the  nitrous  anhydride  inevitably  lost  by  reduction  to 
nitrous  oxide  and  otherwise,  fresh  nitric  acid  is  furnished  by  small 
open  vessels  JV,  containing  sodium  nitrate  and  sulphuric  acid,  placed 
in  the  flues  of  the  pyrite-burners.  About  4  kg.  of  the  nitrate  are 
consumed  for  every  100  kg.  of  sulphur. 

The  immense  size  of  the  chambers  is  necessitated  by  the  fact  that 
the  chemical  action,  although  much  quicker  than  the  direct  oxidation 
of  sulphurous  acid,  is  after  all  rather  slow.  The  presence  of  the  large 
amount  of  atmospheric  nitrogen,  which  diminishes  the  concentration 
of  all  the  interacting  substances,  partly  accounts  for  this  slowness. 
The  acid  which  accumulates  upon  the  floors  contains  but  60  to  70  per 
cent  of  sulphuric  acid,  and  has  a  specific  gravity  of  1.5-1.62.  The 
excess  of  water  is  used  to  facilitate  the  second  action.  It  is  required 
also  in  order  that  the  acid  upon  the  floor  may  not  afterwards  absorb 
and  retain  the  nitrous  anhydride,  for  this  substance  combines  with  an 
acid  containing  more  than  70  per  cent  of  hydrogen  sulphate. 

This  crude  sulphuric  acid  is  applicable  directly  in  some  chemical 
manufactures,  such  as  the  preparation  of  superphosphates  (q.v.)  used 
in  large  amounts  as  a  fertilizer.  In  most  cases,  however,  a  more 
concentrated  sulphuric  acid  is  required.  The  concentration  is 
effected  in  the  first  place  by  evaporation  in  pans  lined  with  lead,  which 
are  frequently  placed  over  the  pyrite-burners  in  order  to  economize 
fuel.  The  evaporation  in  lead  is  carried  on  until  a  specific  gravity 
1.7,  corresponding  to  77  per  cent  concentration,  is  reached.  Up  to 
this  point  the  sulphate  of  lead  formed  by  the  action  of  the  sulphuric 
acid  produces  a  crust  which  protects  the  metal  from  further  action. 
The  insoluble  sulphate  of  lead,  however,  becomes  more  soluble  in 
sulphuric  acid  the  more  concentrated  the  acid  is,  and  the  higher, 
therefore,  its  boiling-point.  When  a  stronger  acid  is  required,  the 
water  is  usually  driven  out  by  heating  the  sulphuric  acid  in  vessels  of 
porcelain  or  platinum,  or  even  of  cast  iron.  Iron  acts  upon  dilute 
sulphuric  acid,  displacing  the  hydrogen-ion,  but  not  upon  the  con- 
centrated acid,  which'  is  not  ionized.  Commercial  sulphuric  acid, 
oil  of  vitriol,  has  a  specific  gravity  1.83-1.84,  and  contains  about  93.5 
per  cent  of  sulphuric  acid. 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         437 

Physical  Properties.  —  Pure  hydrogen  sulphate  (100  per  cent) 
has  a  sp.  gr.  1.85  at  15°.  When  cooled,  it  crystallizes  (m.-p.  10.5°). 
At  150°-180°  the  acid  begins  to  fume,  giving  off  sulphur  trioxide.  At 
330°  it  boils,  but  loses  more  sulphur  trioxide  than  water,  and  finally 
leaves  in  the  retort  an  acid  of  constant  (p.  211)  boiling-point  (338°) 
and  constant  composition  (98.33  per  cent). 

When  hydrogen  sulphate  is  mixed  with  water,  a  considerable  evo- 
lution of  heat  takes  place.  This  heat  of  solution  (p.  203)  receives 
progressively  diminishing  increments  as  more  water  is  added,  until  a 
.very  great  dilution  has  been  reached.  The  total  is  39,170  cal.  This 
heat  of  solution  has  not  been  accounted  for  quantitatively,  but  a  part 
of  it  is  due  to  the  heat  given  out  in  connection  with  the  ionization  of 
the  hydrogen  sulphate.  The  pure  substance,  and  the  concentrated 
acid,  absorb  water  greedily  from  the  moisture  in  the  air  (or  other  gas), 
and  so  are  used  for  drying  gases. 

Impurities.  —  Commercial  sulphuric  acid  is  frequently  brown 
in  color  on  account  of  the  presence  of  fragments  of  straw  which  have 
become  charred  and  finally  completely  disintegrated.  It  contains  also 
lead  sulphate,  which  appears  as  a  precipitate  when  the  acid  is  diluted, 
as  well  as  arsenic  trioxide  and  oxides  of  nitrogen  in  combination,  and 
many  other  foreign  substances  in  small  quantities.  The  pure  sul- 
phuric acid  employed  in  chemical  laboratories  has  received  special 
treatment  for  the  removal  of  these  ingredients.  ' 


Chemical  Properties  of  Hydrogen  Sulphate  HzSO^  —  1.    The 

compound  is  not  exceedingly  stable,  for  dissociation  into  water  and 
sulphur  trioxide  begins  far  below  the  boiling-point  (cf.  p.  437).  The 
vapor  of  the  acid  boiling  at  338°  contains  only  66  per  cent  of  H2S04 
and  34  per  cent  of  H^O  +  SOs,  which  recombine  when  the  vapor  is 
cooled.  The  dissociation  is  practically  complete  at  416°,  as  is  shown 
by  the  density  of  the  vapor.  When  raised  suddenly  to  a  red  heat  it 
is  broken  up  completely  into  water,  sulphur  dioxide,  and  oxygen. 

2.  When  sulphur  trioxide  is  dissolved  in  hydrogen  sulphate,  pyro- 
sulphuric  acid  H^O?,  a  solid  compound,  is  obtained.  Hydrogen 
sulphate  containing  20  per  cent  of  pyro-sulphuric  acid  is  known  as 
"oleum,"  and  is  employed  in  chemical  industries.  The  old  "  fuming," 
or  "  Nordhausen,  "  sulphuric  acid  contained  10-20  per  cent  of  extra 
sulphur  trioxide.  The  salts  of  disulphuric  acid  may  be  made  by 
strongly  heating  the  acid  sulphates,  for  example: 

2NaHS04  <=*  Na2S207  +  H20  t  . 


438  INORGANIC  CHEMISTRY 

In  view  of  this  mode  of  preparation  by  the  aid  of  heat,  they  were 
named  pyrosulphates  (Gk.  wp,  fire).  When  they  are  dissolved  in 
water,  the  acid  sulphates  are  reproduced  by  reversal  of  the  foregoing 
reaction. 

3.  With  salts  which  it  does  not  oxidize,  hydrogen  sulphate  reacts 
by  double  decomposition  and  sets  free  the  corresponding  acid.     The 
actions  are  always  reversible  ones;  but  where  the  new  acid  is  volatile, 
as  in  the  case  of  hydrogen  chloride  (p.  206),  we  are  furnished  with  one 
of  the  cheapest  means  of  preparing  acids. 

Since  hydrogen  sulphate  is  dibasic  (p.  372),  it  forme  both  acid 
and  neutral  salts,  such  as  NaHSO4  and  Na2SO4.  The  acid  sulphates 
are  called  also  bisulphates,  because  they  contain  twice  as  large  a 
proportion  of  SO4  to  metallic  element,  and  require  twice  as  much 
sulphuric  acid  for  their  preparation  as  do  the  neutral  sulphates. 

4.  Sulphuric  acid  combines  vigorously  with  water  to  form  at 
least  one  rather  stable  hydrate,  H^SO^H^O   (m.-p.  8°).      On  this 
account,  sulphuric  acid  is  able  to  take  the  elements  of  water  from 
compounds  containing  hydrogen  and  oxygen,  especially  those  con- 
taining these  elements  in  the  proportion  2H  :  O.     Thus  paper,  which 
is  largely  cellulose  (CeHioOs)*,  wood  which  contains  much  cellulose, 
and  sugar  C^H^On  are  charred  by  it,  and  carbon  is  set  free: 

>  12C  +  11H20. 


The  same  tendency  is  enlisted  to  promote  chemical  actions  in  which 
water  is  formed,  particularly  in  connection  with  the  manufacture  of 
nitroglycerine  (q.v.)  and  guncotton  (q.v.).  For  the  same  reason, 
sulphuric  acid  is  used  in  drying  gases  with  which  it  does  not  interact. 
5.  On  account  of  the  large  quantity  of  oxygen  which  hydrogen 
sulphate  contains,  and  its  instability  when  heated,  it  behaves  as  an 
oxidizing  agent.  This  property  has  already  been  illustrated  in  con- 
nection with  the  action  of  the  acid  upon  carbon,  sulphur,  and  copper 
(p.  425),  hydrogen  sulphide  (p.  416),  zinc  (p.  119),  and,  particularly, 
hydrogen  iodide  (p.  277)  and  hydrogen  bromide  (p.  272).  The 
sulphuric  acid  is  itself  reduced  to  sulphur  dioxide,  and  even  to  free 
sulphur  or  hydrogen  sulphide.  The  metals,  from  the  most  active 
down  to  silver  (p.  404),  are  capable  of  reducing  it,  the  sulphates*  being 
formed.  The  more  active  metals,  like  zinc,  reduce  it  to  hydrogen 

*  Note  that  the  sulphates,  and  not  the  oxides  of  the  metals,  are  produced. 
Oxides  of  metals  could  not  be  formed  in  concentrated  sulphuric  acid,  because 
ihey  interact  with  the  latter  much  more  vigorously  than  do  the  metals,  to  give 
the  sulphates  (cf.  p.  213). 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         439 

sulphide  (p.  416),  the  less  active,  like  copper,  give  sulphur  dioxide 
(p.  425).  Hydrogen  is  not  liberated,  because  no  hydrogen-ion  is 
present  in  concentrated  sulphuric  acid.  Gold  and  platinum  alone  are 
not  attacked,  and  hence  their  use  in  making  sulphuric  acid  stills. 
Free  hydrogen  itself  is  oxidized  to  water  when  passed  into  hydrogen 
sulphate  at  160°:  SO2(OH)2  +  H2  ->  S02  +  2H2O. 

Concentrated  sulphuric  acid  is  used  in  almost  all  chemical  in- 
dustries: for  example,  to  give  sodium  sulphate,  as  a  stage  in  the 
Le  Blanc  process  for  the  manufacture  of  soda;  in  the  refining  of 
petroleum;  .in  the  manufacture  of  fertilizers,  such  as  superphos- 
phate; in  the  preparation  of  nitroglycerine  and  guncotton,  where 
it  assists  the  action  by  removing  water;  and  in  the  production  of 
coal-tar  dyes. 

lonization  of  Dibasic  Acids.  —  An  acid  containing  but  one 
unit  of  hydrogen  in  its  molecule  can  give  but  two  kinds  of  ions. 
Thus,  chloric  acid  gives  only  H+  and  C1O3~.  When  more  than  one 
hydrogen  unit  is  present,  however,  more  than  two  kinds  of  ions  are 
formed.  Thus,  sulphuric  acid,  H2SO4,  produces,  in  the  first  place, 
hydrosulphate-ion  : 


The  latter  is  also  an  acid,  but  is  considerably  less  active  than  sulphuric 
acid.  Hence,  the  further  dissociation  of  this  ion:  HS04~  <=±  H+  + 
SO4=,  lags  considerably  behind  the  primary  dissociation.  In  con- 
centrated solutions  of  the  acid  there  is,  therefore,  much  HSO4~ 
present.  In  very  dilute  solutions,  however,  S04=  predominates.  We 
know  that  HSO4~  is  a  weaker  acid,  and  is  dissociated  with  greater 
difficulty  by  water,  because  acid  salts,  like  KHSO4,  which  give  this 
ion,  are  much  weaker  acids  than  are  acids  like  HC1  and  HC1O3,  with 
which  the  substance  HSO4~  might  fairly  be  compared.  This  be- 
havior is  not  peculiar  to  sulphuric  acid,  but  is  shown  by  all  acids 
containing  more  than  one  hydrogen  unit  in  the  molecule  (cf.  Hydrogen 
sulphide,  p.  417). 

Chemical   Properties   of  Aqueous   Hydrogen   Sulphate.  — 

The  solution  of  sulphuric  acid  is  a  mixture  whose  components  are: 
undissociated  molecules  H2SO4,  hydrogen-ion  H+,  hydrosulphate-ion 
HSO4~~,  and  sulphate-ion  SO4~.  The  chemical  properties  shown  by 
the  solution  are  those  of  one  or  other  of  these  components,  according 
to  circumstances. 


440 


INORGANIC  CHEMISTRY 


Except  in  concentrated  solutions  (normal  or  stronger)  the  oxidizing 
effects  of  the  undissociated,  molecular  substance  are  not  encountered. 
The  temperature  of  the  diluted  acid,  even  when  boiling,  is  not  high 
enough  for  the  purpose.  In  fairly  strong  solutions,  hydrosulphate-ion 
is  plentiful  and  shows  itself  in  the  results  of  electrolysis  (see  p.  449). 

The  presence  of  hydrogen-ion  is  shown  by  all  its  usual  properties 
(p.  373).  In  the  following  table  the  proportion  of  the  whole  of  the 
hydrogen  existing  in  the  form  of  hydrogen-ion  (column  five)  and  its 
concentration  (column  six),  taking  a  normal  solution  of  hydrogen-ion 
containing  1  g.  per  liter  as  standard,  are  shown  (cf.  pp.  182  and  366). 
The  first  three  columns  give  the  concentration  of  the  sulphuric  acid  as 
a  whole,  in  terms  (first  column)  of  the  volume  (liters)  of  liquid  contain- 
ing one  equivalent  (JH2SO4  =  49  g.),  in  terms  (second  column)  of  a 
normal  solution  as  standard,  and  by  per  cent  (third  column),  respec- 
tively. The  fourth  column  shows  the  conductivity  (p.  365). 


V 

H2S04 

PER  CENT 
H2S04. 

K; 

x^Aoo 

H+ 

0.1 

ION 

38.00 

70 

0.18 

1.8AT 

1 

N 

4.79 

198 

0.51 

0.51AT 

10 

o.uv 

0.48 

225 

0.58 

0.058Ar 

100 

0.01  AT 

0.05 

308 

0.79 

0.0079AT 

1000 

O.OOIAT 

0.005 

361 

0.93 

0.00093AT 

oo 

0 

0.00 

388 

1.00 

0.00 

Column  5  thus  states  that  in  a  normal  solution  51  per  cent,  and  in  a 
centi-normal  solution  79  per  cent  of  the  hydrogen  is  ionic. 

Sulphate-ion  SO4=:,  which  is  found  also  in  solutions  of  all  neutral 
and  acid  sulphates,  unites  with  all  positive  ions.  The  product,  when 
insoluble,  appears  as  a  precipitate.  The  introduction  of  barium  ions, 
for  example,  by  adding  a  solution  of  barium  nitrate  or  chloride,  is  em- 
ployed as  a  test  for  sulphate-ion: 

.  Ba++ +  804=4=^8041. 

Since  there  are  other  barium  salts  which  are  insoluble  in  water  (see 
Table  of  solubilities),  but  no  common  ones  which  are  not  decomposed 
by  acids,  dilute  nitric  acid  is  first  added  to  the  solution  supposed  to 
contain  the  sulphate-ion.  The  other  ions,  if  present,  then  give  no 
precipitate  with  barium-ion. 

Dilute  sulphuric  acid  is  used  for  many  purposes.  Thus  it 
forms  the  liquid  in  the  lead  storage  battery,  and  is  employed  for 
cleaning  sheet  iron  before  tinning  and  galvanizing. 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         441 

Sulphates.  —  The   acid   sulphates,  known   also  as  bisulphates 

(p.  438),  may  be  produced  either  by  adding  to  dilute  sulphuric  acid 
half  an  equivalent  of  a  base,  and  evaporating  :  NaOH  -f-  H2SO4  *=> 
H2O  +  NaHSC>4,  or  by  actions  in  which  another  acid  is  displaced  by 
concentrated  sulphuric  acid,  as  in  making  hydrogen  chloride  (p.  206)  . 
These  salts  are  acid  in  reaction,  as  well  as  in  name  (cf.  p.  401),  because 
HS04~,  although  a  weak,  is  not  a  feeble  acid.  When  heated,  they 
yield  pyrosulphates  (p.  438). 

The  normal  (or  neutral)  sulphates  are  obtained  by  complete 
neutralization  and  evaporation,  or  by  the  second  of  the  above  methods 
when  a  sufficient  amount  of  the  salt  and  a  higher  temperature  are  used  : 


NaCl  +  NaHS04  <=±  NajSOi  +  HC1  f  . 

They  are  often  made  also  by  precipitation,  by  oxidation  of  a  sulphide 
at  a  high  temperature,  PbS  -+-  202  —  »  PbSO4,  or  by  addition  of  sul- 
phur trioxide  to  the  oxide  of  a  metal  (p.  431). 

Normal  sulphates  of  many  heavy  metals  decompose  at  a  red  heat, 
those  of  the  trivalent  metals  giving  off  sulphur  trioxide  (p.  430),  and 
those  of  some  bivalent  metals  (e.g.,  Mn,  Ni,  Co),  giving  sulphur 
dioxide  and  oxygen.  The  sulphates  of  potassium,  sodium,  and  others 
of  the  more  active  metals,  and  lead  sulphate,  however,  are  not 
affected  by  heating. 

When  a  sulphate,  or  indeed  any  salt  of  a  sulphur  acid,  is  heated 
strongly  with  carbon,  the  oxygen  is  removed  and  a  sulphide  remains: 
Na^SCX  +  4C  —  »  Na2S  +  4CO.  Upon  this  is  founded  a  general  test 
for  the  presence  of  sulphur  in  any  substance.  The  material  to  be 
tested  is  mixed  with  sodium  carbonate.  A  small  amount  of  the  mix- 
ture is  placed  on  the  end  of  a  match,  which  has  been  charred  and 
rendered  partially  incombustible  by  previous  application  of  sodium 
carbonate.  When  the  end  of  the  match  is  now  held  in  the  reducing 
part  of  the  Bunsen  flame,  the  compound  of  sulphur,  if  it  contains 
oxygen,  is  reduced  to  the  form  of  sulphide.  This,  by  interaction  with 
the  carbonate,  gives  sodium  sulphide,  Na^S.  When  the  product  of  the 
reduction  is  placed  upon  a  silver  coin  and  moistened,  the  sodium  sul- 
phide, if  present,  produces  a  black  stain  of  silver  sulphide.  This  is 
known  as  the  hepar  test,  hepar  being  an  old  name  for  a  sulphide. 

Constitution  of  Hydrogen  Sulphate.  —  The  formula  which 


we  assign  to  sulphur  trioxide  is  0  =  S  ^    .     It  is  in  general  our 


442  INORGANIC  CHEMISTRY 

desire  to  use  the  smallest  possible  valence,  but  here  no  reduction  can 
be  effected  below  the  value  6  for  the  sulphur,  unless  we  join  the 

/O 
oxygen  units  to  one  another,  as  in  the  formula  O  =  S  x  I  .     This, 

O-H 

however,  would  suggest  a  relationship  to  hydrogen  peroxide,    i          , 

O  —  H 

which  is  not  confirmed,  for  hydrogen  peroxide  cannot  be  made  from 
sulphuric  acid.  Assuming,  therefore,  the  above  formula  for  sulphur 
trioxide,  the  addition  of  the  elements  of  water  to  it  in  the  simplest 
fashion  results  in  the  structures: 


or  i  .. 

O'?*0  H-0/S^0 

Jl 

The  second  of  these  two  modes  of  disposing  of  the  water  is  the  one 
which,  in  parallel  cases,  is  usually  most  feasible.  Hardly  any  alter- 
native to  it  is  possible,  for  example,  in  representing  the  action  in 
which  quicklime  is  slaked  (p.  149)  : 


This  represents  the  change  with  little  derangement  of  the  original 
structure  and  without  alteration  in  the  valence,  while  the  first  unwar- 
rantably increases  the  valence  to  ten.  There  are  other  objections  to 
the  first  formula.  In  it  the  hydrogen  is  supposed  to  unite  more 
immediately  with  the  sulphur,  whereas,  when  the  free  elements  are 
concerned,  hydrogen  actually  combines  more  readily  with  oxygen,  and 
forms  a  more  stable  compound  with  it  than  with  sulphur.  Again, 
compounds  like  hydrogen  sulphide,  H  —  S  —  H,  in  which  the  hydro- 
gen is  undoubtedly  united  to  sulphur,  are  but  slightly  ionized,  and  are 
feeble  acids,  while  hydrogen  sulphate  is  highly  ionized. 

Another  fact  is  more  satisfactorily  accounted  for  by  the  second 
formula.  The  addition  of  chlorine  to  sulphur  dioxide  must  be  shown 
thus: 

,0  Cl  O 

S^     +Cfc->       )S' 

^o  cr     ^o 

for  chlorine  has  a  much  greater  tendency  to  unite  with  sulphur  than 
with  oxygen.  When  the  product,  sulphuryl  chloride,  is  brought  iu 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         443 

contact  with  water,  sulphuric  acid  and  hydrogen  chloride  are  pro- 
duced. Since  water  has  the  formula  H  —  O  —  H,  and  two  molecules 
of  water  are  used,  this  action  is  most  simply  accounted  for,  with  the 
minimum  of  disturbance  in  both  molecules,  by  imagining  the  opera- 
tion to  take  place  as  follows: 


H-0- 
H-0- 


H    OI 


H    Cl 


:  \  <2 


The  hydrogen  chloride  is  eliminated,  and  the  other  units  of  hydrogen, 
originally  without  doubt  attached  to  oxygen  in  the  water,  may  be  pre- 
sumed to  be  still  connected  with  that  oxygen  when  they  enter  the 
molecule  of  hydrogen  sulphate. 

This  illustration  shows  the  sort  of  reasoning,  based  upon  the 
chemical  properties  and  the  modes  of  formation  of  a  substance,  which 
lead  us  to  the  devising  of  an  appropriate  graphic  or  structural  formula 
(cf.  p.  322).  The  latter  is  not  supposed  to  represent  the  actual 
physical  structure  of  the  molecule,  but  simply  to  be  a  diagrammatic 
representation  of  the  chemical  relations  of  the  constituents  and  of  the 
chemical  behavior  of  the  whole.  Formulae  of  this  kind  are  in  con- 
tinual use  in  the  study  of  the  compounds  of  carbon,  but  are  seldom 
required  outside  of  that  region. 

OTHER  ACIDS  OF  SULPHUR 

Hyposulphurous  Acid  H2S2O4.  —  The  zinc  salt  of  this  acid 
crystallizes  out  when  zinc  dust  acts  upon  a  solution  of  sulphur  dioxide 
in  absolute  alcohol: 

Zn-f-2SO2-*ZnS204. 

Moissan  found  that  when  sulphur  dioxide  was  led  over  sodium  hy- 
dride, sodium  hyposulphite  was  formed :  2NaH  +  2SO2  — >  Na2S2C>4 
+  H2. 

Commercially,  a  solution  containing  the  sodium  salt  is  made  by 
the  interaction  of  zinc  with  a  solution  of  sodium  bisulphite  charged 
with  excess  of  sulphur  dioxide: 

2NaHS03  +  S02  +  Zn  -»  Na2S2O4  +  ZnS03  +  H20. 

The  salts  are  rapidly  oxidized  by  the  air,  giving  sulphites  and  then 
sulphates.  The  above  solution  of  sodium  hyposulphite  is  used  in 
indigo  dyeing,  and  with  other  vat  dyes,  on  account  of  its  high  reducing 


444  INORGANIC  CHEMISTRY 

power.     Indigo  Ci6Hi0N202,  which  is  insoluble,  is  reduced  by  the  salt 
to  indigo-white  C^H^N^  which  passes  into  solution: 

Na2S2O4  +  2H20  -»  2NaHS03  (+  2H)  (1) 

C16H10N202  (+  2H)  ->  C16H12N202  _     (2) 
2H2O  +  Ci6H10N2O2  -»  2NaHSO3  +  Ci6H]2N202 


When  cloth  saturated  with  the  mixture,  however,  is  exposed  to  the 
air,  the  indigo-white  undergoes  oxidation,  and  blue,  insoluble  indigo 
is  formed  once  more  (see  Dyeing). 

The  acid  is  formed  when  sulphurous  acid  surrounds  the  negative 
electrode  in  an  electrolytic  cell  : 

2H2SO3  +  H2  ->  H2S204  +  2H2O, 

and  was  first  named  (after  the  discoverer)  Schtitzenberger's  acid. 
The  sodium  salt  is  formed  in  the  same  way  from  sodium  bisulphite. 

Sulphurous  Acid.  —  This  term  is  applied  to  the  solution  of 
sulphur  dioxide  in  water.  A  portion  of  the  sulphur  dioxide  remains 
dissolved  physically,  while  another  portion  is  in  combination  with  the 
water,  forming  sulphurous  acid.  This  in  turn  is  ionized,  and  chiefly, 
after  the  manner  of  the  weaker  dibasic  acids,  into  two  ions,  H+  and 
HSO3~.  A  little  SO3=  is  formed  from  the  latter.  There  are  thus  in 
such  a  solution  four  mutually  dependent  equilibria: 


S02  (gas)  ^±S02  (dslvd)  +  H20<=»H2S03  <=»  H 

When  the  solution  is  heated,  uncombined  sulphur  dioxide  is  dis- 
engaged as  a  gas.  The  equilibria  being  thus  disturbed,  the  ions  of  the 
.acid  unite,  the  acid  molecules  decompose,  and  soon  all  the  above 
actions  are  completely  reversed  and  the  whole  of  the  gas  passes  off. 
Conversely,  when  a  base  furnishing  hydroxide  ions  is  added  to  the 
solution  of  the  acid,  the  hydrogen  ions  disappear,  forming  water,  and 
the  above  actions  all  proceed  in  a  forward  direction  until,  with  a  half- 
equivalent  of  the  base,  the  whole  of  the  material  has  been  converted 
into  the  form  HSO3~,  in  association,  of  course,  with  the  positive  ions 
of  the  base.  With  a  full  equivalent,  neutralization  follows  and  SO3~ 
is  the  product. 

Properties  of  Sulphurous  Acid.  —  The  acid  is  so  unstable  that 
it  cannot  be  obtained  excepting  in  solution  in  water.  Chemically  it  is 
a  comparatively  weak  acid. 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         445 

As  a  reducing  agent,  it  is  slowly  oxidized  by  free  oxygen,  turning 
into  sulphuric  acid.  Sugar  and  glycerine  act  as  negative  contact 
agents  and  make  the  oxidation  much  slower.  It  is  rapidly  oxidized 
by  active  oxidizing  agents.  Thus,  when  free  halogens  are  added  to 
the  solution,  sulphuric  acid  and  the  hydrogen  halide  are  formed: 

H2S03  +  H20  +  I2  +±  H2S04  +  2HI. 

In  the  particular  case  of  iodine  this  action  takes  place  only  in  very 
dilute  solution,  since  concentrated  sulphuric  acid  decomposes  hydro- 
gen iodide  (cf.  p.  277)  and  the  action  is  reversed.  This  interaction  is 
used  in  chemical  analysis  as  a  means  of  estimating  the  quantity  of 
sulphurous  acid  in  a  liquid  (cf.  p.  277). 

Hydrogen  peroxide,  potassium  permanganate,  and  other  oxidizing 
agents  convert  the  substance  into  sulphuric  acid  likewise.  It  should 
be  noted  that  in  these  oxidations  we  have,  not  an  addition  of  oxygen 
to  SO2,  but  to  the  SO3~  or  HSO3~  ion  of  the  acid,  whereby  it  passes  into 
the  SO4=  ion  of  sulphuric  acid.  The  ion  is  much  more  easily  oxidized 
than  is  free  sulphur  dioxide  itself. 

When  heated  alone,  in  a  sealed  glass  tube  (150°),  the  acid  reduces 
part  of  itself  to  sulphur,  and  a  part  is  oxidized  to  sulphuric  acid: 

3H2S03  ->  2H2S04  +  H20  +  S. 

Sulphurous  acid  has  the  power  of  uniting  directly  with  many 
organic  coloring  matters  and,  since  the  products  of  this  union  are 
usually  colorless,  it  is  employed  as  a  bleaching  agent.  It  is  especially 
useful  with  materials  like  silk,  wool,  and  straw,  which  are  likely  to  be 
destroyed  by  hypochlorous  acid.  Sunlight  causes  the  dissociation  of 
these  colorless  compounds,  and  so,  with  use,  straw  hats  slowly  recover 
their  original  color.  As  a  disinfectant  it  acts  by  addition  likewise. 

As  a  dibasic  acid,  sulphurous  acid  forms  normal  salts  like  Na2SO3, 
and  acid  salts  like  NaHSO3. 

Consecutive  Reactions.  —  There  are  many  chemical  reactions 
that  proceed  in  two  stages,  which  can  be  carried  out  separately. 
This  is  the  case  with  the  two  reactions  used  in  the  chamber  process 
(p.  433).  The  actions  are  consecutive  (p.  272),  because  the  second 
uses  materials  produced  by  the  first.  It  may  be  noted  that,  if  the 
second  action  is  as  speedy  as  the  first,  or  speedier,  then  no  inter- 
mediate products  will  be  detectable.  This  is  the  case  with  the 
chamber  process  reactions,  when  sufficient  steam  is  introduced, 
for  under  these  circumstances  no  solid  nitrosylsulphuric  acid  is 


446  .  INORGANIC   CHEMISTRY 

deposited.  If  the  second  reaction  is  slower  than  the  first,  then  the 
products  of  the  first  reaction  will  accumulate,  and  become  notice- 
able. 

The  conception  of  consecutive  reactions  enables  us  to  under- 
stand and  remember  some  facts.  For  example,  it  was  mentioned 
that  when  dry  sulphur  is  oxidized,  we  obtain  sulphur  dioxide,  but 
when  moist  sulphur  is  oxidized,  by  the  air  or  otherwise,  the  only 
product  is  sulphuric  acid  (p.  414).  This  change  may  be  conceived 
of  as  proceeding  in  two  stages: 

S  +  02  +  H20-*H2S03, 
2H2SO3  +  O2    ->  2H2SO4, 

which  would  be  consecutive  reactions.  Since  oxidation  of  solid 
sulphur  can  proceed  only  on  the  surface,  it  is  slow.  Since  the  sul- 
phurous acid  is  dissolved,  and  every  molecule  of  it  is  accessible 
to  the  dissolved  oxygen,  or  oxidizing  agent,  the  second  action  should 
be  speedier  and  consume  the  product  of  the  first  action  as  fast  as  it 
is  formed.  It  is,  therefore,  quite  natural  that  no  sulphurous  acid 
should  be  detectable  when  water  (or  its  vapor)  is  present. 

Illustration  of  the  Effect  of  Concentration  on  Speed  of  In- 
teraction. —  The  oxidation  of  sulphurous  acid  by  iodic  acid  HI03  may 
be  used  to  show  [Lect.  exp.]  the  effect  of  concentration  on  the  speed 
of  an  action  (p.  291).  The  iodic  acid  may  most  readily  be  made  by 
dissolving  potassium  iodate  KI03  and  sulphuric  acid  together  in 
water,  in  such  quantities  as  would  give  a  N/2  solution  of  each.  When 
1  c.c.  of  this  N/2  iodic  acid  is  added  to  100  c.c.  of  filtered  starch 
emulsion,  and  the  whole  is  mixed  with  an  equal  volume  of  water 
containing  1  c.c.  of  N/2  sulphurous  acid,  the  blue  color  produced  by 
the  liberated  iodine  appears  suddenly  after  the  lapse  of  a  minute  or 
more: 

2HI03  +  5H2S03  -»  5H2S04  +  H2O  +  I2. 

With  double  the  above  quantities  in  the  same  amount  of  water,  that  is, 
with  double  concentrations,  the  speed  of  the  action  is  greatly  increased 
and  the  iodine  becomes  visible  in  less  than  half  the  time. 

Sulphites.  —  The  acid  sulphites  of  the  alkali  metals  (i.e.,  of 
potassium  and  sodium)  are  acid  in  reaction,  owing  to  the  appreciable 
dissociation  of  the  ion  HSO3~.  The  acid  being  a  weak  one,  however, 
solutions  of  the  normal  salts,  Na2SO3,  etc.,  are  alkaline  towards  litmus 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         447 

(p.  399).  The  sulphites  are  readily  decomposed  by  acids  giving  free 
sulphurous  acid,  and  the  latter  decomposes,  yielding  sulphur  dioxide 
(p.  424). 

Calcium  bisulphite  solution,  Ca(HSO3)2,  is  used  to  dissolve  the 
lignin  out  of  wood  chips  employed  in  the  manufacture  of  paper. 
About  30  per  cent  of  the  wood  is  lignin.  The  rest  is  cellulose 
(CeHioOs)^  and  constitutes  the  prepared  pulp. 

When  heated,  sulphites  undergo  decomposition.  The  sulphates, 
being  the  most  stable  of  all  the  salts  of  sulphur  acids,  are  formed  when 
the  salts  of  any  of  those  acids  are  decomposed  by  heating.  The  nature 
of  the  particular  salt  determines  what  other  products  shall  appear. 
Here,  one.  molecule  of  the  sulphite  furnishes  three  atoms  of  oxygen, 
sufficient  to  oxidize  three  other  molecules,  and  leaves  one  molecule  of 
sodium  sulphide  behind  (compare  effect  of  heating  sulphurous  acid, 
alone,  p.  445)  : 

4Na2SO3  -»  Na2S  +  3Na2SO4. 

The  acid  sulphites  (bisulphites)  first  lose  sulphurous  acid,  before 
changing  in  this  way.  Thus,  sodium-hydrogen  sulphite  begins  by 
decomposing  as  follows: 

2NaHSO3  r*  Na2S03  +  H2S03  (or  H20  +  S02) 


and  at  a  higher  temperature  the  sulphite  NasSOs  decomposes  as 
explained  above. 

The  acid  salts  of  volatile  acids,  when  heated,  all  decompose  in  this 
way  (cf.  pp.  281,  447). 

The  sulphites  are  as  readily  oxidized  as  is  the  acid  itself.  They  are 
slowly  converted,  both  in  solution  and  in  the  solid  form,  by  the  influ- 
ence of  the  oxygen  of  the  air,  into  sulphates.  It  is  interesting  to  note 
that  the  addition  of  sugar  or  glycerine  to  a  solution  of  a  sulphite 
reduces  the  speed  of  oxidation  by  free  oxygen  very  markedly.  These 
substances  act  as  contact  agents;  and  the  present  case  shows  that 
agents  of  this  kind  may  not  only  increase  the  speed  of  actions,  which 
is  their  usual  function,  but  may  also  have  a  restraining  influence. 

Thiosulphuric  Acid  H2S2O3.  —  This  acid  is  not  known  in  the 
free  condition,  but  its  salts  are  in  common  use  in  the  laboratory  and 
commercially.  Sodium  thiosulphate,  for  example,  is  prepared  by 
boiling  a  solution  of  sodium  sulphite  with  free  sulphur.  The  action 
is  something  like  the  addition  of  oxygen  to  sulphurous  acid: 

+  S  -»  Na&Os    or    SO3=  +  S  ->  S203=. 


448  INORGANIC   CHEMISTRY 

The  product,  thiosulphate  of  sodium,  is  used  in  photography  as  a 
solvent  for  salts  of  silver  (fixing  bath),  and  is  commonly  (but  in- 
correctly) called  "  hypo." 

By  the  addition  of  acids  to  a  solution  of  sodium  thiosulphate,  the 
thiosulphuric  acid  is  set  free,  but  the  latter  instantly  decomposes, 
giving  a  precipitate  of  sulphur: 

Na2S203  +  2HC1  <±  H2S203  +  2NaCl, 


Even  carbon  dioxide  from  the  air,  giving  carbonic  acid  (q.v.),  pro- 
duces this  effect  slowly  in  fixing  solutions.  The  actions  being  revers- 
ible, preliminary  addition  of  a  sulphite  to  the  solution  helps  to  sustain 
the  reverse  action,  in  which  sulphurous  acid  is  a  factor,  and  so  pre- 
serves the  solution.  The  delay  in  the  appearance  of  the  precipitate 
of  sulphur  in  dilute  solutions  is  due  to  the  temporary  existence  of  a 
supersaturated  solution  (cf.  p.  193)  of  the  free  element.  This  is  shown 
by  the  fact  that  instant  neutralization  of  the  free  acid  does  not 
prevent  the  ultimate  appearance  of  the  sulphur. 

Iodine  acts  upon  sodium  thiosulphate  solution,  giving  sodium 
tetrathionate: 

I2  -»  2NaI 


This  action  is  used,  by  employment  of  a  standard  solution  of  sodium 
thiosulphate,  for  estimating  quantities  of  free  iodine  in  analysis.  The 
disappearance  of  the  color  of  the  latter  indicates  that  a  sufficient 
amount  of  the  salt  has  been  employed.  When  chlorine-water  (p.  223) 
is  used,  the  oxidation  is  more  complete.  The  products  are  sodium 
sulphate,,  sulphuric  acid,  and  hydrochloric  acid  : 

Na*S203  +  4HC10  +  H2O  ->  Na*S04  +  H2SO4  +  4HC1. 


In  consequence  of  the  very  great  amount  of  free  chlorine  which  the 
sodium  thiosulphate  is  thus  able  to  transform,  it  is  employed,  as 
antichlor,  for  the  purpose  of  removing  chlorine  from  bleached  fabrics 

Persulphuric  Acid  H2S2O8.  —  The  salts  of  this  acid  are  coming 
into  use  for  commercial  purposes  and  for  "  reducing"  negatives  in 
photography.  When  a  discharge  of  electricity  is  passed  through  a 
mixture  of  sulphur  trioxide  and  oxygen,  drops  of  liquid  are  formed 
which  appear  to  have  the  composition  S2O7,  and  when  dissolved  in 
water  give  dilute  persulphuric  acid,  S207  +  H20  —  »  H2S2O8.  More 
significant  of  its  relations  is  its  formation,  to  some  extent,  when 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         449 

concentrated  sulphuric  acid  and  a  strong  solution  of  hydrogen  per- 
oxide are  mixed: 

2H2S04  +  H202  +±  H2S208  +  2H20. 

Tnis  action  is  reversible.*  Under  some  circumstances,  particularly 
by  using  pure  (100  per  cent)  hydrogen  peroxide  and  sulphur  trioxide, 
a  monobasic  acid,  monopersulphuric  acid  (Caro's  acid,  m.-p.  45°)  is 
formed:  H2SO4  +  H2O2  <=±  H2SO5  +  H2O.  Interesting  in  its  way, 
also,  is  its  production  in  the  electrolysis  of  aqueous  sulphuric  acid: 


This  action  is  most  conspicuous  in  rather  concentrated  (50  per  cent) 
solutions,  in  which  hydrosulphate-ion  is  plentiful  (cf.  p.  439),  and 
when  a  small  anode,  resulting  in  severe  crowding  of  the  HS04  radicals 
as  they  are  liberated,  is  employed.  The  salts  were  first  prepared  by 
electrolyzing  ammonium  or  sodium-hydrogen  sulphate  NaHS04  in 
concentrated  solution  (Hugh  Marshall).  The  persulphuric  acid, 
formed  by  the  union  of  the  negative  ions  in  pairs,  undergoes  double 
decomposition  with  the  excess  of  sodium  bisulphate,  and  the  less 
soluble  sodium  persulphate  Na2S2O8  crystallizes  out.  The  other  salts 
are  made  by  double  decomposition  from  this  one. 

The  persulphates  decompose  readily  when  heated,  yielding  pyro- 
sulphates  and  oxygen: 

2K2S208-^2K2S2O7  +  O2. 

The  solution  of  the  acid  is  an  active  oxidizing  agent: 
H2S2O8  +  H2O  ->  2H2S04  (+  O). 

Polythionic  Acids.  —  Di-,  tri-,  tetra-,  and  pentathionic  acid  (p. 
431)  are  all  formed  simultaneously  (along  with  free  sulphur)  when 
sulphur  dioxide  and  hydrogen  sulphide  gases  are  passed  alternately 

*  This  action  and  the  next  are  not  classifiable  under  any  of  the  ten  kinds 
formerly  discussed  (p.  228).  They  consist  in  the  union  of  H  and  OH  to  form 
water: 

HSO4j  H  +  HO  :  OH  +  H  j  SO4H  -»  2H2O  +  (SO4H)2. 

Neutralizations  (p.  387)  they  are  not,  because  the  interacting  substances  are 
both  acids.  Just  as  the  loss  of  water  from  one  acid  gives  an  anhydride,  so  here, 
the  loss  of  water  between  two  acids  gives  a  mixed  anhydride  (see  Chlorosul- 
phuric  acid,  below). 


450  INORGANIC  CHEMISTRY 

into  water,  although  the  gases  themselves  (p.  416)  interact  to  produce 
simply  free  sulphur  and  water: 

H2S  +  3S02->H2S406, 
2H2S  +  6S02  ->  H2S306  +  H2S606, 
3H2S  +  9S02-»H2S206  +  2H2S506. 

Most  of  these  acids  and  their  salts  are  of  minor  interest  and  need  not 
be  discussed. 

The  production  of  sodium  tetrathionate  by  the  action  of  iodine 
upon  sodium  thiosulphate  has  already  been  mentioned  (p.  448). 

When  manganese  dioxide  is  treated  with  cold  sulphurous  acid,  it 
interacts  rapidly  and  a  solution  of  manganous  dithionate  is  obtained  : 

Mn02  +  2H2SO3  ->  MnS2O6  +  2H20. 

The  salts  of  these  acids  are  in  many  cases  fairly  stable,  but  the 
acids  themselves  decompose  readily  when  set  free. 

COMPOUNDS  OF  SULPHUR  WITH  CHLORINE  AND 
FLUORINE 

Sulphur  Monochloride  S2C^.  —  When  chlorine  gas  is  passed 
over  heated  sulphur  it  is  absorbed,  and  a  reddish-yellow  liquid, 
boiling  at  138°,  is  obtained.  The  molecular  weight  of  this  substance, 
as  shown  by  the  density  of  its  vapor,  indicates  that  it  possesses  the 
formula  S2C12.  When  thrown  into  water  it  is  rapidly  hydrolyzed, 
producing  sulphur  dioxide  and  free  sulphur: 

2H20  -»  SO2  +  4HC1  +  3S. 


Sulphur  itself  dissolves  very  freely  in  the  monochloride.  The 
monochloride  is  employed  in  vulcanizing  rubber. 

By  surrounding  sulphur  monochloride  with  a  freezing  mixture, 
and  treating  it  with  excess  of  chlorine,  a  liquid  dichloride  SC12,  and  a 
tetrachloride  SCU  can  be  formed.  Both  are  unstable.  Moissan  pre- 
pared sulphur  hexafluoride  SFe,  which  is  a  gas  at  room  temperature 
(b.-p.  —50°).  Unlike  the  chloride,  it  is  not  hydrolyzed  by  water. 

Thionyl  Chloride  SOCl*.  —  By  the  action  of  sulphur  dioxide 
gas  upon  phosphorus  pentachloride,  part  of  the  oxygen  in  the  former 
is  replaced  by  chlorine: 

S02  -|-  PC16  -»  SOC12  +  POC13. 
The  products  are  thionyl  chloride  and  phosphorus  oxychloride.     The 


THE  OXIDES  AND  OXYGEN  ACIDS  OF  SULPHUR         451 

former  is  a  colorless  liquid,  boiling  at  78°,  and  is  separated  from  the 
latter  (b.-p.  107°)  by  fractional  distillation  (see  Petroleum).  It  is  de- 
composed immediately  on  contact  with  water: 

SOC12  +  2H2O  ->  H2SO3  +  2HC1. 

Sulphuryl  Chloride  SO2O2.  —  Sulphur  dioxide  and  chlorine 
gases  unite  when  exposed  to  direct  sunlight  to  form  a  liquid  known  as 
sulphuryl  chloride  S02C12.  When  camphor  is  introduced  into  the 
vessel  the  union  takes  place  much  more  rapidly,  owing  to  some  cata- 
lytic effect  of  this  substance.  The  compound  is  a  colorless  liquid, 
boiling  at  69°.  With  water  it  gives  sulphuric  acid  and  hydrogen 
chloride  (p.  443)  .  When  a  strictly  limited  amount  of  water  is  supplied 
a  partial  action  of  the  same  nature  occurs,  and  the  product  is  known 
as  chlorosulphuric  acid: 

/  PI 
S02C12  +  H20  ->  S02  +  HCL 


This  intermediate  compound  may  be  formed  also  by  the  addition  of 
hydrogen  chloride  to  sulphur  trioxide. 

Exercises.  —  1.   What  ground  is  there  for  assigning  the  formula 
SO2  instead  of  S204  to  sulphur  dioxide  (p.  427)? 

2.  Explain  why  nitric  acid  is  completely  displaced  by  the  action 
of  sulphuric  acid  on  sodium  nitrate  (p.  436). 

3.  How  many  times,  on  an  average,  does  a  molecule  of  nitrous 
anhydride  go  through  the  cycle  of  changes  by  which  sulphuric  acid  is 
produced  before  it  is  eliminated  in  some  other  form  (p.  436)? 

4.  Make  a  list  of,  and  classify,  the  various  applications  of  sul- 
phuric acid  to  the  liberation  of  other  acids. 

5.  Formulate  the  behavior  of  the  hydrosulphate-ion   (p.  439) 
when  a  solution  of  barium  chloride  is  added  to  a  rather  concentrated 
solution  of  sulphuric  acid. 

6.  Can  you  give  any  reasons  for  preferring  to  regard  KHSCX,  and 
substances  like  it,  as  acid  salts  rather  than  double  salts  of  the  form 
K2S04,H2S04  (p.  401)? 

7.  From  a  consideration  of  the  facts  given  in  the  text,  and  of  the 
physical  conditions,  account  for  the  fact  that  hyposulphurous  acid  is 
oxidized  by  free  oxygen  first  to  sulphurous  acid  and  then  to  sulphuric 
acid  (p.  443),  while  moist  free  sulphur,  even  with  oxidizing  agents, 
gives  sulphuric  acid  directly  (p.  445)? 


452  INORGANIC  CHEMISTRY 

8.  Why  were  not  pyrosulphuric  acid  and  monopersulphuric  acid 
placed  in  the  list  on  p.  431  (cf.  p.  487)? 

9.  Write  in  ionic  form  the  equation  for  the  interaction  of  sodium 
thiosulphate  and  iodine  in  aqueous  solution. 

10.  Restate  the  import  of  the  following  sentence  in  terms  of 
experimental  facts:    Liquefied  sulphur  dioxide  "ionizes  substances 
dissolved  in  it  as  well  as  does  water"  (p.  427). 

11.  What  are  the  relative  volumes,  (a)  of  sulphur  dioxide  and 
nitrogen  (p.  9)  resulting  trom  the  roasting  of  pyrite  (p.  424),  (6) 
of  air  and  sulphur  dioxide  in  making  sulphuric  acid? 

12.  Assign  to  the  proper  classes  of  ionic  actions  (pp.  402-406),  (a) 
the  action  of  iodine  on  sulphurous  acid  (p.  445),  (6)  of  sulphur  on 
sodium  sulphite  (p.  447),  (c)  the  ways  of  forming  persulphuric  acid 
(pp.  448-449). 


CHAPTER  XXII 

SELENIUM  AND   TELLURIUM.     THE   CLASSIFICATION   OF 
THE    ELEMENTS 

ALONG  with  sulphur,  chemists  group  two  other  elements,  selenium 
(Se,  at.  wt.  79.2)  and  tellurium  (Te,  at.  wt.  127.5).  If  the  nature  of 
the  chief  compounds  of  sulphur  is  kept  in  mind,  the  close  analogy 
between  the  nature  and  chemical  behavior  of  the  three  elements  and 
their  corresponding  compounds  will  be  noticed  at  once  (see  Chemical 
relations  of  the  sulphur  family,  below). 

SELENIUM  SE 

Occurrence  and  Properties  of  the  Element.  —  Selenium  (Gk. 
<r€\Tqjnrj,  the  moon)  occurs  free  in  some  specimens  of  native  sulphur, 
and  in  combination  often  takes  the  place  of  a  small  part  of  the  sulphur 
in  pyrite  Fe&2  (Berzelius,  1817).  It  is  found  free  in  the  dust-flues  of 
the  pyrite-burners  of  sulphuric  acid  works.  The  familiar  forms  are, 
the  red  precipitated  variety,  which  is  amorphous  and  soluble  in 
carbon  disulphide,  and  the  lead-gray,  semi-metallic  variety,  obtained 
by  slow  cooling  of  melted  selenium,  which  is  insoluble,  and  melts  at 
217°.  In  the  latter  form  it  has  some  capacity  for  conducting  electri  > 
ity,  which  is  greatly  increased  by  exposure  to  light  in  proportion  to 
the  intensity  of  the  illumination.  A  photometer,  using  this  property, 
has  been  devised  by  Joel  Stebbins  (1914),  for  measuring  the  rela- 
tive intensity  of  the  light  of  different  stars.  Selenium  boils  at  688°, 
and  at  high  temperatures  has  a  vapor  density  corresponding  to  the 
formula  862. 

The  element  unites  directly  with  many  metals,  burns  in  oxygen  to 
form  selenium  dioxide,  and  unites  vigorously  with  chlorine.  It  is 
used  in  glass-making. 

Hydrogen  Selenide.  —  Ferrous  selenide,  made  by  heating  iron 
filings  with  selenium,  when  treated  with  concentrated  hydrochloric 
acid  gives  hydrogen  selenide: 

FeSe  +  2HC1  <±  H2Se  f  +  FeCl2. 
453 


454  INORGANIC  CHEMISTRY 

The  compound  is  a  poisonous  gas,  which  possesses  an  odor  recalling 
rotten  horse-radish,  and  is  soluble  in  water.  The  solution  is  faintly 
acid  in  reaction,  and  deposits  selenium  when  exposed  to  the  action  of 
the  air  (cf.  p.  417).  The  gas  dissociates  when  heated  (cf.  p.  415). 
Other  selenides,  which,  with  the  exception  of  those  of  potassium  and 
sodium,  are  insoluble  in  water,  may  be  precipitated  by  leading  the 
gas  into  solutions  of  soluble  salts  of  appropriate  metals  (cf.  p.  421). 


Selenium  Dioxide  and  Selenious  Acid.  —  The  dioxide 
is  a  solid  body  formed  by  burning  selenium  or  evaporating  a  solution 
of  selenious  acid  H2SeO3.  The  latter  may  be  made  by  dissolving 
the  dioxide  in  hot  water,  or  by  oxidizing  selenium  with  boiling  nitric 
acid  or  aqua  regia  (q.v.).  Unlike  sulphur  (p.  414),  the  element  gives 
little  of  the  higher  acid  H2Se04  by  this  treatment.  In  aqueous 
solution  the  acid  is  easily  reduced,  even  by  sulphurous  acid,  to 
selenium: 

H2Se03  +  2H2S03  -»  2H2SO4  +  H2O  +  Se. 

Selenic  Acid.  —  No  trioxide  is  known.  Selenic  acid  H2SeO4,  a 
white  solid,  is  made  in  solution  by  oxidizing  silver  selenite  with 
bromine-water  (which  contains  hypobromous  acid,  cf.  p.  223)  and 
filtering: 

Br2  +  H2O  <F±  HBr  +  HBrO 
Ag2SeO3  +  HBrO  ->  Ag2Se04  +  HBr 
2HBr  +  Ag2SeQ4  ->  2AgBr  |  +  H2SeO4 
Br2  +  H2O  +  Ag2SeO3  ->  2AgBr  J  +  H2SeO4 

It  is  itself  a  powerful  oxidizing  agent  and,  even  in  dilute  solu- 
tion, liberates  chlorine  from  hydrochloric  acid:  H2Se04  +  2HC1  —  •> 
H2SeO3  +  H20  +  C12.  Sulphuric  acid  (cf.  p.  438),  on  the  other  hand, 
is  an  oxidizing  agent  only  in  somewhat  concentrated  form,  and  even 
then  it  can  oxidize  hydrobromic  acid  (p.  272),  but  not  hydrochloric 
acid. 

Chlorides  of  Selenium.  —  The  chlorides  are  formed  by  direct 
union  of  the  elements.  The  tetrachloride,  a  yellow  crystalline  sub- 
stance, is  formed  when  the  monochloride  is  heated  :  2Se2Cl2  —  ->  3Se  + 
SeCl4,  the  behavior  in  the  case  of  sulphur  chlorides  being  just  the 
inverse  of  this  (p.  450). 


SELENIUM  AND  TELLURIUM  455 


TELLURIUM  TE 

Tellurium  (Lat.  tellus,  the  earth)  occurs  in  sylvanite  in  combina- 
tion with  gold  and  silver.  It  is  a  white,  metallic,  crystalline  sub- 
stance, melting  at  452°  (b.-p.  1400°).  When  formed  by  precipitation 
it  is  a  black  powder.  It  conducts  electricity  to  some  extent.  The 
vapor  density  corresponds  to  the  formula  Tea.  The  free  element 
unites  with  metals  directly,  and  burns  in  air  to  form  the  dioxide. 

Hydrogen  telluride  H2Te  is  made  by  the  action  of  acids  on 
metallic  tellurides,  and  its  aqueous  solution  is  rapidly  oxidized  by  air 
with  precipitation  of  tellurium.  The  tellurides  of  the  alkali  metals 
are  soluble  in  water,  the  others  are  insoluble. 

Tellurious  acid  H2Te03  is  formed  by  oxidizing  the  element  with 
nitric  acid,  and  is  a  crystalline  solid,  little  soluble  in  water.  It  is  a 
feeble  acid,  of  which  many  salts  have  been  made.  It  is  also  some- 
what basic,  a  sulphate  2Te02,SO3  and  a  nitrate  Te2O3(OH)NO3  being 
known.  In  this  respect  it  differs  markedly  from  sulphurous  acid. 

Telluric  acid  is  made  by  oxidizing  tellurious  acid  in  aqueous  solu- 
tion with  chromic  acid  (p.  418).  It  is  difficultly  soluble  in  water.  It 
does  not  affect  indicators,  and  is  therefore  actually  more  feebly  acidic 
than  is  hydrogen  sulphide.  It  is  obtained  as  a  solid  having  the  com- 
position HeTeOe  (or  3H2O,Te03)  on  evaporating  the  solution.  When 
heated,  this  body  loses  water,  some  of  the  trioxide  TeOs  being 
formed.  The  last  is  a  yellow  solid,  which  shows  no  tendency  to 
recombine  with  water,  in  this  respect  resembling  silica  (q.v.).  Tel- 
lurates  of  the  alkali  metals  may  be  made  by  heating  the  tellurites 
with  potassium  or  sodium  nitrate:  K2Te03  +  KNO3  — >  K2Te04  + 
KN02. 

Tellurium  forms  two  very  stable  chlorides,  TeCl2  and  TeCU,  which 
are  decomposed  by  water.  The  second,  however,  exists  in  solution 
with  excess  of  hydrogen  chloride:  TeCl4  +  3H2O^±H2Te03  +  4HC1, 
showing  the  tellurious  acid  to  be  basic  in  properties  and  the  element 
tellurium  to  be,  to  a  certain  degree,  a  metallic  element  (see  Chap. 
XXXII). 

We  have  now  a  means  of  deciding  whether  a  substance  containing  the 
elements  of  water,  such  as,  for  example,  telluric  acid,  should  be  written,  H6TeO6, 
or  as  a  hydrate  H2TeO4,2H2O.  Baker  and  Adlam  have  shown  that  plates  made 
of  true  hydrates  are  permeable  by  water  vapor,  but  that  crystals  of  anhydrous 
substances,  such  as  cupric  sulphate,  are  not  so  permeable.  They  found  that 
crystals  of  telluric  acid  are  not  permeable  by  water  vapor,  so  that  HeTeOe  is  the 
correct  formula. 


456  INORGANIC  CHEMISTRY 

The  Chemical  Relations  of  the  Sulphur  Family.  —  It  will 
be  seen  that  sulphur,  selenium,  and  tellurium  are  bivalent  elements 
when  combined  with  hydrogen  or  metals.  In  combination  with 
oxygen  they  form  unsaturated  compounds  of  the  form  XIVC>2,  while 
their  highest  valence  is  found  in  SOa,  TeOs,  and  H2Se04,  where  they 
must  be  sexivalent.  The  general  behavior  of  corresponding  com- 
pounds is  very  similar.  At  the  same  time,  there  is  in  all  cases  a  pro- 
gressive change  as  we  proceed  from  sulphur  through  selenium  to 
tellurium.  The  elementary  substances  themselves,  for  example, 
become  more  like  metals,  physically,  and  they  show  higher  and 
higher  melting-points.  The  affinity  for  hydrogen  decreases,  as  is 
shown  by  the  increasing  ease  with  which  the  compounds  H2X  are 
oxidized  in  air.  The  affinity  for  oxygen  likewise  decreases,  for  the 
elements  become  increasingly  difficult  to  raise  to  the  highest  state 
of  oxidation.  On  the  other  hand,  the  tendency  to  form  higher  chlo- 
rides becomes  greater.  We  note  also  that  the  compounds  H2X04 
become  less  arid  less  active  as  acids,  and  a  distinct  basic  tendency 
begins  to  assert  itself. 

THE  PERIODIC  SYSTEM 

Classification,  or  the  arrangement  of  facts  on  the  basis  of  likeness, 
is  part  of  the  method  of  science.  In  chemistry  the  multitude  of  facts 
is  not  less  than  in  other  sciences,  and  the  necessity  of  arrangement 
equally  urgent.  It  is  needed  to  make  possible  the  systematic  descrip- 
tion of  the  ascertained  facts,  and  to  furnish  a  guide  in  investigation,  by 
suggesting  stochastic  hypotheses  (p.  176),  and  so  pointing  out  direc- 
tions in  which  new  facts  of  interest  may  be  found.  Thus,  as  an  aid  to 
memory,  we  have  treated  the  halogens  as  one  family  and  S,  Se,  and 
Te  as  another.  In  each  case,  we  have  presented  the  properties 
common  to  all  members  of  the  group,  and  have  then  pointed  out  the 
differences.  Again,  in  investigation,  as  soon  as  we  have  discovered 
that  sulphur  and  selenium  are  allied  elements,  we  realize  the  direction 
in  which  fruitful  results  may  be  expected,  and  we  proceed  to  make 
the  corresponding  compounds  and  to  note  the  resemblances  and 
differences  in  the  conditions  for  preparation  and  in  the  properties 
of  the  compounds  obtained. 

At  first  sight,  the  most  definite  method  of  classification  would 
appear  to  be  the  grouping  of  elements  of  like  valence.  But  this  brings 
together  sodium  and  chlorine  —  an  element  whose  hydrogen  com- 
pound is  unstable  and  without  markedly  characteristic  properties, 


THE  PERIODIC  SYSTEM  457 

and  whose  hydroxyl  compound  is  an  active  base,  with  an  element 
whose  hydrogen  compound  is  an  active  acid  and  whose  hydroxyl 
compound  is,  in  a  feeble  degree,  an  acid  also.  This  method  homolo- 
gates similar  and  contrasting  elements  indiscriminately. 

Metallic  and  Non- Metallic  Elements.  —  Thus  far  we  have 
found  the  division  into  metallic  and  non-metallic  elements  very 
serviceable  for  classification  in  terms  of  chemical  relations  (p.  150). 
This  distinction  we  shall  continue  to  employ.  The  metallic,  or 
positive  elements  (p.  357),  (1)  form  positive  radicals  and  ions  contain- 
ing no  other  element  (cf.  p.  375).  Thus  the  metals  give  sulphates, 
nitrates,  carbonates,  and  other  salts,  which  furnish  a  metallic  ion, 
such  as  Na+  or  K+,  together  with  the  ions  S04=,  N03~,  and  CO3=.  (2) 
Their  hydroxides,  KOH,  Ca(OH)2,  etc.,  give  the  same  metallic  ion, 
and  the  rest  of  the  molecule  forms  hydroxide-ion.  That  is  to  say, 
their  hydroxides  are  bases  and  their  oxides  are  basic.  The  metallic 
elements  often  enter,  but  only  with  other  elements,  into  the  composi- 
tion of  a  negative  ion,  as  is  the  case  with  manganese  in  K.MnO4, 
with  chromium  in  K2.Cr2O7,  and  with  silver  in  K.Ag(CN)2.*  But 
the  most  definitely  metallic  elements  form  with  oxygen  such  a  nega- 
tive ion  only  while  exhibiting  a  different  valence  from  that  which  they 
possess  when  acting  as  positive  elements.  Thus,  manganese  when 
a  positive  element  has  the  valences  two  and  three,  MnO  and  Mn2O3, 
Mn.Cl2  and  Mn.Cl3,  Mn.SC>4  and  Mn2.(SO4)3,  etc.,  while  in  perman- 
ganates we  have  potentially  the  oxide  Mn2Oy,  (K2O,Mn2O7  = 
2KMnO4),  in  which  it  is  septivalent.  The  graphic  formulae  express 
the  difference  in  valence: 

ci  xci  ^o 

Mn(  Mn-Cl  K  -  O  -  Mn=0 

Cl  XC1  ^O 

If  we  knew  only  the  compounds  in  which  manganese  is  septivalent, 
we  should  regard  it  as  a  non-metallic  element  pure  and  simple,  the 
metallic  appearance  of  the  free  element  to  the  contrary  notwith- 
standing. 

The  non-metallic  or  negative  elements,  (1)  are  found  chiefly  in 
negative  radicals  and  ions.  They  form  no  nitrates,  sulphates,  carbon- 
ates, etc.,  for  they  could  not  do  so  without  themselves  alone  con- 
stituting the  positive  ion.  We  have  no  such  salts  of  oxygen,  sulphur, 

*  The  mode  of  division  into  ions  is  shown  by  the  position  of  the  period  in  the 
formula. 


458  INORGANIC  CHEMISTRY 

carbon,  or  phosphorus,  for  example.  (2)  Their  hydroxides,  although 
their  formulae  may  be  written  C1O2OH,  P(OH)3,  SO2(OH)2,  furnish 
no  hydroxide  ions,  as  this  would  involve  the  same  consequence. 
These  hydroxides  are  divided  by  dissociation,  in  fact,  so  that  the 
non-metal  forms  part  of  a  compound  negative  radical,  and  the  other 
ion  is  hydrogen-ion,  C1O3.H,  P03H.H2,  SO4.H2.  (3)  Their  halogen 
compounds,  like  PCla  (p.  210)  and  S2C12  (p.  450),  are  completely 
hydrolyzed  by  water,  and  the  actions  are  not,  in  general,  reversible. 
The  halides  of  the  typical  metals  are  not  hydrolyzed  (see  Chap. 
XXXII),  and  with  those  that  are  not  typical,  the  action  is  reversible. 
The  distinctions  are  not  perfectly  sharp,  however.  Thus,  zinc 
(q.v.)  gives  both  salts  like  the  sulphate  Zn.SCX  and  chloride  Zn.Cl2, 
and  compounds  like  sodium  zincate  (p.  122)  ZnO2H.Na,  showing  the 
same  valence  in  both  classes: 

Cl  O  -  Na 

Zn'  Zn' 

Cl  X0-H 

Its  hydroxide  ionizes  in  two  ways,  Zn.  (OH)2  and  ZnO2H.H.  Similarly 
tellurious  acid  H2TeO3  acts  both  as  acid  and  base  (p.  455).  We  shall 
find  this  double  behavior  conspicuous  in  the  compounds  of  arsenic  and 
antimony.  In  spite  of  the  partial  merging  of  the  two  classes  of  ele- 
ments, however,  the  general  distinction  is  worth  preserving  (see  Chap. 
XXXII). 

Classification  by  Atomic  Weights.  —  A  closer  discrimination 
than  that  furnished  by  these  two  categories  is  required,  however,  and 
a  study  of  the  order  into  which  the  elements  fall  when  arranged 
according  to  their  atomic  weights  has  provided  this  to  some  extent. 

The  first  indication  of  a  significant  relation  between  the  atomic 
weights  and  the  properties  of  the  elements  was  given  by  a  fact  noted 
by  Dobereiner  (1829).  He  drew  attention  to  the  existence  of  closely 
similar  elements  in  sets  of  three  (triads),  where  the  central  element 
was  intermediate  in  properties  between  the  two  others,  and  the  atomic 
weight  of  the  central  element  was  almost  the  exact  arithmetical  mean 
of  the  weights  of  the  other  two.  The  three  following  triads  illustrate 
this  relation: 

Chlorine   ....  35.46  Sulphur  ....     32.06  Calcium  ....     40.07 

Bromine    ....  79.92  Selenium     ...     79.2  Strontium   ...     87.63 

Iodine 126.92  Tellurium    .    .    .    127.5  Barium    .    .    .    .    137.37 

Mean  of  Cl  and  I,  81.2  Mean  of  S  and  Te,  79 . 8  Mean  of  Ca  and  Ba,  88 . 7 


THE  PERIODIC  SYSTEM  459 

Newlands  (1863-4)  discovered  a  surprising  regularity  when  the 
elements  were  placed  in  the  order  of  ascending  atomic  weight.  Omit- 
ting hydrogen  (at.  wt.  1)  the  first  seven  are:  lithium  (7),  glucinum 
(9),  boron  (11),  carbon  (12),  nitrogen  (14),  oxygen  (16),  fluorine  (19). 
These  are  all  of-  totally  different  classes,  and  include  first  a  metal 
forming  a  strongly  basic  hydroxide,  then  a  metal  of  the  less  active 
sort,  then  five  non-metals  of  increasingly  negative  character,  the 
last  being  the  most  active  non-metal  known.  The  next  element  after 
fluorine  (19)  is  sodium  (23),  which  brings  us  back  sharply  to  the 
elements  that  form  strongly  basic  hydroxides.  Omitting  none,  the 
next  seven  elements  are:  sodium  (23),  magnesium  (24.3),  aluminium 
(27),  silicon  (28.3),  phosphorus  (31),  sulphur  (32),  chlorine  (35.5). 
In  this  series  there  are  three  metals  of  diminishing  positiveness, 
followed  by  four  non-metals  of  increasing  negative  activity,  the  last 
being  a  halogen  very  like  fluorine.  On  account  of  the  fact  that  each 
element  resembles  most  closely  the  eighth  element  beyond  or  before  it 
in  the  list,  the  relation  was  called  the  law  of  octaves.  After  chlorine 
the  octaves  become  less  easy  to  trace.  Potassium  (39)  follows  chlo- 
rine and  corresponds  satisfactorily  to  sodium,  but  it  is  not  until  seven- 
teen successive  elements  have  been  set  down  that  we  reach  one  closely 
resembling  chlorine,  namely,  bromine. 

That  this  periodicity  in  chemical  nature  is  more  than  a  coincidence 
is  shown  by  the  fact  that  the  valence  and  even  the  physical  properties, 
such  as  the  density,  show  a  similar  fluctuation  in  each  series,  and 
recurrence  in  the  following  one.  In  the  first  two  series  the  compounds 
with  other  elements  are  of  the  types: 

LiCl,    GlCla,    BC13,    CCU,   NH3,  OH2,  FH 
Li2O,    G1O,      B2O3,    CO2,    N205 

NaCl,  MgCl2,  A1C13,  SiCl4,  PH3,   SH2,    C1H 
Na^O,  MgO,    A12O3,  SiO2,    P205,  SO3,    C12O7 

Thus  the  valence  towards  chlorine  or  hydrogen  ascends  to  four  and 
then  reverts  to  one  in  each  octave.  The  highest  valence,  shown  in 
oxygen  compounds,  ascends  from  lithium  to  nitrogen  with  values  one 
to  five,  and  then  fails  because  compounds  are  lacking.  In  the  second 
octave,  however,  it  goes  up  continuously  from  one  to  seven. 

Again,  the  densities  of  the  elements  in  the  second  series,  using 
the  data  for  red  phosphorus  and  liquid  chlorine,  are:  % 

Na  0.97,  Mg  1.75,  Al  2.67,  Si  2.49,  P  2.14,  S  2.06,  Cl  1.33. 


460  INORGANIC  CHEMISTRY 

Of  greater  significance  chemically  are  the  related  numbers  represent- 
ing the  volumes  in  cubic  centimeters  occupied  by  a  gram-atomic 
weight  of  each  element  (the  atomic  volumes) : 

Na  24,  Mg  14,  Al  10,  Si  11,  P  14,  S  16,  Cl  27. 

A  similar  regular  fluctuation  is  shown  by  all  the  physical  properties  of 
corresponding  compounds. 

Mendelejeff 's  Scheme.  —  In  1869  Mendelejeff  published  an 
important  contribution  towards  adjusting  the  difficulty  which  the  ele- 
ments following  chlorine  presented,  and  developed  the  whole  concep- 
tion so  completely  that  the  resulting  system  of  classification  has  been 
connected  with  his  name  ever  since.  Almost  simultaneously  Lothar 
Meyer  made  similar  suggestions,  but  did  not  urge  them  with  the  same 
conviction  or  elaborate  them  so  fully.  The  following  table,  in  which 
the  atomic  weights  are  expressed  in  round  numbers,  'is  a  modification 
of  one  of  Mendelejeff 's. 

The  chief  change  from  the  arrangement  in  simple  octaves  is  that 
the  third  series,  beginning  with  potassium,  is  made  to  furnish  material 
for  two  octaves,  potassium  to  manganese  and  copper  to  bromine, 
and  is  called  a  long  series.  The  valences  fall  in  with  this  plan  fairly 
well.  Copper,  while  usually  bivalent,  forms  also  a  series  of  com- 
pounds in  which  it  appears  to  be  univalent.  Iron,  cobalt,  and  nickel 
cannot  be  accommodated  in  either  octave,  as  their  valences  are 
always  two  or  three,  so  these  elements  are  set  off  in  a  column  by 
themselves.  At  the  time  Mendelejeff  made  the  table,  three  places  in 
the  third  (long)  series  had  to  be  left  blank,  as  a  trivalent  element  [Sc] 
was  lacking  in  the  first  octave,  and  a  trivalent  [Ga]  and  a  quadri- 
valent one  [Ge]  in  the  second.  These  places  have  since  been  filled, 
as  we  shall  presently  see.  The  first  two  (the  short)  series  have  been 
split  in  the  table,  as  lithium  and  sodium  closely  resemble  potassium, 
while  the  remaining  members  of  these  series  fall  more  naturally  over 
the  corresponding  elements  of  the  second  octave  of  the  third  series. 

The  fourth  series  (long)  is  nearly  complete.  It  begins  with  an 
active  alkali  metal,  rubidium,  and  ends  with  iodine,  a  halogen. 
The  rule  of  valence  is  strictly  preserved  throughout  the  series,  and  in 
general  the  elements  fall  below  those  which  they  most  closely 
resemble. 

The  fifth,  sixth,  and  seventh  (long)  series  are  incomplete,  but  the 
order  of  the  atomic  weights  and  the  valence  enables  us  satisfactorily 
to  place  many  of  the  elements.  The  chemical  relations  to  elements 


THE  PERIODIC  SYSTEM 


461 


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462  INORGANIC  CHEMISTRY 

of  the  fourth  series  justify  the  position  assigned  to  each.  Caesium, 
for  example,  is  the  most  active  of  the  alkali  metals;  barium  has 
always  been  classed  with  strontium,  and  bismuth  with  antimony. 
There  are,  however,  about  twelve  more  or  less  rare  elements  which 
cannot  easily  be  fitted  into  the  vacant  places. 

In  two  cases  a  slight  displacement  of  the  order  according  to  atomic 
weights  was  necessary.  Cobalt  was  put  before  nickel  because  it 
resembled  iron  more  closely,  although  the  atomic  weights  are  in  the 
opposite  order.  Tellurium  and  iodine  are  placed  in  that  order  to 
bring  them  into  the  sulphur  and  halogen  groups,  respectively.  Their 
valences  and  other  chemical  relations  both  require  this.  The  gen- 
eral agreement,  however,  is  very  remarkable. 

General  Relations  in  the  System.  —  In  every  octave  the 
valence  towards  oxygen  ascends  from  one  to  seven,  while  that  towards 
hydrogen,  in  the  cases  of  the  four  last  elements  (when  they  combine 
with  hydrogen  at  all),  descends  from  four  to  one.  The  atomic 
volume  disregards  the  subsidiary  octaves  in  the  long  series.  It 
descends  towards  the  middle  of  each  series  (long  or  short),  and 
ascends  again  towards  its  initial  value.  The  other  physical  properties 
fluctuate  within  the  limits  of  each  series  in  a  similar  way.  The  values 
of  each  physical  constant  for  corresponding  members  of  the  successive 
series  do  not  exactly  coincide,  however.  A  progressive  change, 
as  we  descend  each  vertical  column,  is  the  rule.  Thus  the  densities 
of  the  alkali  metals  rise  from  lithium  (0.53)  to  caesium  (1.87).  In 
the  same  group  the  melting-points  descend  from  lithium  (186°)  to 
caesium  (26.5°). 

It  must  be  stated  that  no  mathematical  (quantitative)  relation  be- 
tween the  values  for  any  property  and  the  values  of  the  atomic  weights 
has  been  discovered;  only  a  general  (quantitative)  relationship  can 
be  traced.  Anticipating  the  discovery  of  some  more  exact  mode 
of  stating  the  relationship  in  each  case,  and  remembering  that  similar 
values  of  each  property  recur  periodically,  usually  at  intervals  corre- 
sponding to  the  length  of  an  octave  or  series,  the  principle  which  is 
assumed  to  underlie  the  whole,  the  periodic  law,  is  stated  thus: 
All  the  properties  of  the  elements  are  periodic  functions  of  their 
atomic  weights. 

That  the  chemical  relations  of  the  elements  vary  just  as  do  the 
physical  properties  of  the  simple  substances  is  easily  shown.  Thus, 
each  series  begins  with  an  active  metallic  (positive)  element,  and  ends 
with  an  active  non-metallic  (negative)  element,  the  intervening 


THE  PERIODIC  SYSTEM  463 

elements  showing  a  more  or  less  continuous  variation  between  these 
limits.  Again,  the  elements  at  the  top  are  the  least  metallic  of  their 
respective  columns.  As  we  descend,  the  members  of  each  group  are 
more  markedly  metallic  (in  the  first  columns),  or,  what  is  the  same 
thing,  less  markedly  non-metallic  (in  the  later  columns;  cf.  p.  455). 

In  the  first  series  boron  is  the  first  non-metallic  element  we 
encounter.  In  the  second  series  silicon  is  the  first  such  element.  In 
the  third  there  is  more  difficulty  in  deciding.  Titanium,  vanadium, 
and  germanium  are  usually,  though  with  questionable  propriety, 
classed  as  metallic  elements.*  Selenium  is  undoubtedly  a  non- 
metallic  element.  Arsenic  is,  on  the  whole,  a  non-metallic  element. 
In  the  fourth  series  tellurium  is  commonly  considered  to  be  the 
first  non-metallic  element.  Thus  a  zigzag  line,  shown  in  the  table, 
separates  all  the  non-metallic  elements  from  the  rest  of  the  elements, 
and  confines  them  in  the  right-hand  upper  corner. 

A  more  compact  form  of  the  table,  also  based  upon  one  of  Mendele- 
jeff's,  is  printed  at  the  end  of  this  book,  opposite  the  rear  cover. 
The  only  difference  between  this  and  the  other  is  that  the  two  octaves 
of  each  long  series  have  been  placed  in  the  same  set  of  seven  main 
columns.  The  irregular  sets  of  three  elements,  consisting  of  the 
iron,  palladium,  and  platinum  groups,  occupy  a  column  on  the  right 
of  the  main  columns,  and  are  often  called  collectively  the  eighth 
group.  The  recently  discovered  noble  gases,  found  chiefly  in  the 
air,  have  been  placed  at  the  left-hand  side.  Since  they  do  not  enter 
into  combination  at  all,  their  valence  may  appropriately  be  given 
as  zero.  With  the  exception  of  argon,  the  values  of  their  atomic 
weights  agree  well  with  this  assignment.  Hydrogen  is  the  only 
common  element  whose  place  is  still  in  debate.  Many  rare  elements 
have  also  been  omitted,  and  tantalum  has  been  placed  immediately 
after  cerium.  The  valence  is  shown  by  the  general  formula  at  the 
head  of  each  column. 

Chemical  Relations  and  the  Periodic  System.  —  It  is  so 

easy  to  confuse  the  qualities  which  enable  us  to  place  an  element 
in  a  suitable  group  or  family,  with  those  which  are  of  no  service 
in  this  respect,  although  they  may  be  used  to  justify  the  assign- 
ment after  it  has  been  made,  that  it  seems  best  to  give  a  concrete 
illustration,  with  the  halogen  family  as  the  example.  In  the  follow- 

*  In  discussing  chemical  relations,  the  term  metallic  element  is  preferable 
to  metal.  The  free  element  (e.g.,  arsenic)  may  have  the  luster  of  a  metal,  and 
yet  the  element,  in  combination,  may  be  non-metallic  or  negative, 


464 


INORGANIC  CHEMISTRY 


ing  table  there  are  named:  (1)  a  few  specific  physical  properties  of 
a  simple  uncombined  substance,  (2)  a  few  chemical  properties  of  a 
simple  uncombined  substance,  (3)  the  chemical  relations  of  the 
element  as  judged  by  its  properties  when  in  combination,  or  by 
the  behavior  of  its  compounds: 


Phys.  Props, 
of  Simple 

Chem.  Props. 
Substance 

Chem.  Relations  of  Element. 

Color 
Density 
Crit.  temp. 
M.-p.,  b.-p. 

Molar  wt. 
Combines  with 

Atomic  Wt. 
Valence 
Metallic  or  non-metallic  element 

Displaces 

Loses  these  after  combination 

Possesses  these  in  combination 

Could  not  group  elements  by 
these  properties 

Do  group  the  element  by  these 
relations 

Might  oxygen  fairly  be  included  in  the  halogen  family?  It  is  almost 
colorless,  so  it  might  be  the  first  member  of  the  family,  iodine,  the 
most  strongly  colored,  being  the  last.  Its  density  is  lower  than  that 
of  any  halogen,  as  are  also  its  critical  temperature,  boiling-point, 
and  melting-point.  We  cannot  decide  on  the  basis  of  physical 
properties  alone.  Or,  using  the  properties  in  the  second  column, 
can  we  ascertain  whether  sulphur  might  fairly  be  considered  a  halo- 
gen? Its  molar  weight  is  256,  higher  than  that  of  iodine,  although 
like  iodine,  it  gives  lower  values  at  higher  temperatures,  so  it  might 
be  the  last  member  of  the  family.  It  combines  directly  with  prac- 
tically all  the  metals,  just  as  do  the  halogens.  All  the  halogens 
displace  it  from  combination  with  metals.  It  unites  directly,  and 
vigorously  with  oxygen,  while  iodine  forms  more  stable  compounds 
with  oxygen  than  do  bromine  or  chlorine.  The  last  two  -facts  would 
then  place  sulphur  as  last  member  of  the  series.  But,  should  it  be 
in  this  family  at  all?  The  chemical  properties  of  the  simple  sub- 
stances do  not  enable  us  to  decide. 

When  we  examine  the  column  of  chemical  relations,  however, 
we  find  oxygen  to  be  bivalent,  while  the  halogens  are  univalent 
(and  septivalent  towards  oxygen),  so  oxygen  does  not  belong  to  this 
group.  Again,  sulphur  is  non-metallic  (oxides  acidic,  halides  hydro- 
lyzecl  completely),  just  like  the  halogens.  But  sulphur  is  bivalent 
or  sexivalent,  and  its  atomic  weight  places  it  between  fluorine  and 
chlorine,  and  not  after  iodine,  §o  that  it  cannot  be  included  in  this 


THE  PERIODIC  SYSTEM  465 

gs'oup.  Thus,  it  is  evident  that  we  decide  which  elements  form  a 
coherent  family  by  consideration  of  the  chemical  relations,  i.e.,  by 
the  qualities  each  element  shows  while  in  combination. 

Applications  of  the  Periodic  System.  —  The  system  has  found 
application  chiefly  in  four  ways : 

1.  In  the  prediction  of  new  elements.     Mendelejeff  (1871)  drew 
attention  to  the  blank  then  existing  between  calcium  (40)  and  titanium 
(48).     He  predicted  that  an  element  to  fit  this  place  would  have  an 
atomic  weight  44  and  would  be  trivalent.     From  the  nature  of  the 
surrounding  elements,  he  very  cleverly  deduced  many  of  the  physical 
and  chemical  properties  of  the  unknown  element  and  of  its  com- 
pounds.    He  named  it  eka-boron  (Skr.  eka,  one).     In  1879  Nilson 
discovered  scandium  (44),  and  its  behavior  corresponded  closely  with 
that  predicted  for  eka-boron.     Mendelejeff  also  described  accurately 
two  other  elements,  likewise  unknown  at  the  time.     They  were  to 
occupy  vacant  places  between  zinc  and  arsenic,  and  were  named 
eka-aluminium  and  eka-silicon.     In  1875  Lecoque  de  Boisbaudran 
found  gallium,  and  in  1888  Winkler  discovered  germanium,  and  these 
blanks  were  filled.     Other  possible  elements,  such  as  eka-manganese, 
were  described,  but  still  remain  to  be  discovered. 

2.  By  enabling  us  to  decide  on  correct  values  for  the   atomic 
weights  of  some  elements,  when  the  equivalent  weights  have  been 
measured,  but  no  volatile  compound  is  known   (cf.  p.  63,  239). 
Thus  the  atomic  weight  of  uranium  was  thought  to  be  120  until  it  was 
observed  that  no  place  near  to  antimony  (120)  remained  unoccupied. 
With  the  value  240  (now  238.2),  the  element  was  accommodated  at 
the  foot  of  the  column  containing  those  which  it  most  resembled. 
Again,  the  equivalent  weight  of  indium  was  38  and,  as  the  element  was 
supposed  to  be  bivalent,  it  received  the  atomic  weight  76.     It  was 
quite  out  of  place  near  arsenic  (75),  however,  being  decidedly  a  me- 
tallic element.     As  a  trivalent  element  with  the  atomic  weight  115, 
it  fell  between  cadmium  and  tin.     Later  work  fully  justified  the 
change.     Still  again,  glucinum,  with  equivalent  weight  4.5,  resembled 
aluminium  so  strongly  that  it  was  thought  to  be  trivalent,  like  that 
element,  and  to  have  the  atomic  weight  13.5.     But  the  only  vacancy 
in  the  first  series  then  existing  was  between  lithium  (7)  and  boron  (11) 
and  subsequent  investigation  showed  that  the.  properties  of  glucinum 
placed  it  most  fittingly  in  that  position  as  a  bivalent  metallic  element 
with  the  atomic  weight  9.     Recently,  radium  (q.v.)  has  been  dis- 
covered, and  found  to  have  the  equivalent  weight  113  and  to  resemble 


466  INORGANIC  CHEMISTRY 

barium.  If,  like  barium,  it  is  bivalent,  it  occupies  a  place  under 
this  element,  in  the  last  series.  The  atomic  weights  of  cobalt  and 
nickel,  and  of  tellurium  and  iodine,  however,  cannot  be  adjusted.^ 

3.  By  suggesting  problems  for  investigation.     The  periodic  sys- 
tem has  been  of  constant  service  in  the  course  of  inorganic  research, 
and  has  often  furnished  the  original  stimulus  to  such  work  as  well. 
For  example,  the  atomic  weights  of  the  platinum  metals  at  first  placed 
them  in  the  order,  Ir.(197),  Pt  (198),  Os  (199),  although  the  resem- 
blance of  osmium  to  iron  and  ruthenium  would  have  led  us  to  expect 
that  this  element  should  come  first.     For  similar  reasons  platinum 
should  have  come  last,  under  palladium.     A  reinvestigation  of  the 
atomic  weights,  suggested  by  these  considerations,  was  undertaken 
by  Seubert,  and  the  old  values  were  found  in  fact  to  be  very  in- 
accurate.    He  obtained:  Os  =  191,  Ir  =  193,  Pt  =  195. 

Again,  the  atomic  weight  of  tellurium  bore  the  value  128,  when  the 
table  was  first  constructed,  and  it  was  confidently  expected  that  re- 
examination  would  bring  this  value  below  that  of  iodine  (then  127, 
now  126.92).  Several  most  careful  studies  of  the  subject  have  been 
made,  using  different  methods.  It  seems  probable  that  the  real  value 
of  the  atomic  weight  is  not  far  from  Te  =  127.5,  and  therefore  more 
than  half  a  unit  greater  than  that  of  iodine.  Since,  however,  quanti- 
tative correspondence  is  found  nowhere  in  the  system,  the  existence 
of  marked  inconsistencies  like  this  need  not  shake  our  confidence  in 
its  value  when  it  is  used  with  due  consideration  of  the  degree  of 
correspondence  to  be  expected. 

Originally  lead,  although  it  fell  in  the  fourth  column,  possessed 
only  one  compound,  Pb02,  in  which  it  seemed  to  be  undoubtedly  quad- 
rivalent. Search  for  salts  of  the  same  form,  however,  speedily  yielded 
the  tetrachloride  PbCU,  tetracetate,  and  many  others.  The  existence 
of  osmic  acid  OsO4,  and  a  corresponding  compound  of  ruthenium, 
suggests  that  other  compounds  of  the  elements  of  the  eighth  group, 
displaying  the  valence  eight,  may  be  capable  of  preparation.  The 
collocation  of  copper,  silver,  and  gold,  in  the  same  column  with  the 
alkali  metals,  is  not  at  present  perfectly  satisfactory,  and  suggests 
the  advisability  of  strengthening  their  position,  if  possible,  by  further 
investigation. 

In  the  same  way,  incorrect  values  of  many  physical  properties 
have  been  detected,  .and  have  been  rectified  by  more  careful 
work. 

4.  By  furnishing  a  comprehensive  classification  of  the  elements, 
arranging  them  so  as  to  exhibit  the  relationships  among  the  physical 


THE  PERIODIC  SYSTEM  467 

and  chemical  properties  of  the  elements  themselves  and  of  their 
compounds.  Constant  use  will  be  made  of  this  property  of  the 
table  in  the  succeeding  chapters.  Having  disposed  of  the  halogen 
and  sulphur  families  (excepting  the  oxygen  compounds  of  the  former), 
situated,  respectively,  in  the  seventh  and  sixth  columns  of  the  table 
(at  the  end  of  this  book),  we  shall  presently  take  up  nitrogen  and 
phosphorus  from  the  right  side  of  the  fifth  column.  Then  from  the 
fourth  column,  we  shall  select  carbon  and  silicon,  and  from  the 
third  boron,  leaving  the  other  more  decidedly  metallic  elements  for 
later  treatment. 

Inadequacy,  of  the  Periodic  System.  —  The  periodic  system 
is  often  described  as  if  it  furnished  a  classification  of  the  properties 
of  chemical  substances  which  was  complete  in  its  scope,  and  ideal 
in  its  exactness.  This,  however,  is  far  from  being  the  case. 

The  order  of  activity  (E.  M.  Series)  of  the  metals  (pp.  129,  404) 
and  of  the  non-metals  (p.  284)  summarizes  many  properties,  and 
explains  many  features  of  the  chemical  behavior  of  the  elements. 
This  list  is  scattered  through  the  periodic  table  (compare  both), 
without  any  trace  of  regularity.  The  order  of  solubility  of  the 
sulphides,  a  valuable  property  in  qualitative  analysis,  is  practically 
coincident  with  the  E.  M.  series  (see  under  Cadmium),  and  is  not 
to  be  found  in  the  periodic  system. 

The  periodic  system  concentrates  attention  too  largely  on  one 
of  the  valences  of  each  element.  Thus,  for  manganese,  it  focuses 
attention  on  the  septivalent  form  in  the  permanganates.  But  man- 
ganous  salts  are  more  like  the  ferrous,  the  cobaltous,  the  chromous, 
and  other  sets  of  salts,  none  of  which  are  in  the  same  column  of  the 
table.  Similarly,  the  manganic  salts  are  like  the  ferric  salts  and 
the  salts  of  aluminium.  Again,  copper  is  univalent  in  one  series  of 
salts,  but  in  its  better  known  salts  it  is  bivalent.  Silver,  which 
belongs  to  the  same  periodic  family  is  always  univalent,  while  gold, 
also  in  the  same  family,  is  univalent  or  trivalent,  and  in  the  latter 
case  is  almost  wholly  a  non-metallic  element.  Indium  has  three 
sets  of  salts,  and  similar  remarks  could  be  made  about  it  and  about 
the  very  many  other  multivalent  elements.  If  it  were  possible  to 
place  each  element  in  several  different  columns,  one  for  each  of  the 
valences  that  it  shows,  the  table  would  then  include  far  more  of  the 
properties  of  the  elements.  But  this  cannot  be  done,  for,  according 
to  the  order  of  magnitude  of  the  atomic  weights,  there  is  but  one 
place  for  each  element.  In  other  words,  the  periodic  system  largely 


468  INORGANIC  CHEMISTRY 

ignores  the  variety  of  different  classes  of  chemical  relations  which  an 
element  with  several  valences  always  shows. 

Attention  has  also  been  called  to  about  twelve  rare  elements  for 
which  suitable  positions  cannot  be  found  amongst  the  vacant  places 
in  the  table.  Then,  also,  the  physical  properties  of  elements  and 
compounds  follow  the  order  of  the  elements  in  a  family  in  a  general 
way.  But  no  formula  can  be  devised  by  which  the  properties  can  be 
calculated  from  the  atomic  weight  and  other  independent  data. 
Often,  one  or  two  physical  properties  do  not  even  follow  the  order 
of  the  elements  in  the  family.  The  values  of  the  physical  properties, 
the  number  and  kinds  of  valences,  and  so  forth,  could  be  predicted 
from  the  table  only  in  a  general  way,  and  sometimes  the  prediction 
would  be  entirely  false.  The  curve  obtained  when  the  atomic 
volumes  are  plotted  against  the  atomic  weights  follows  the  different 
series  with  surprising  regularity,  but  atomic  volumes  cannot  be  used 
'n  analysis,  or  in  most  kinds  of  chemical  work,  and  therefore  lack 
practical  value. 

These  remarks  are  made,  merely  to  guard  against  the  supposition 
that  anything  like  a  complete  account  of  chemistry  can  be  extracted 
from  the  periodic  system.  This  system  can  be  used,  and  we  shall 
use  it  where  it  is  of  value,  but  we  must  not  allow  the  science  to  be 
distorted  or  emasculated  by  ignoring  the  important  relations  of 
which  the  system  takes  no  notice. 

THE  ATOMIC  NUMBERS 

When  the  water  of  hydration  escapes  as  vapor  from  a  hydrate, 
the  crystals  frequently  crumble  to  dust  (p.  154).  There  are,  however, 
interesting  exceptions.  Minerals  of  the  zeolite  class,  for  example, 
can  be  deprived  of  water  without  losing  their  coherence.  When  the 
desiccated  specimen  is  placed  in  contact  with  water  vapor  at  a  suf- 
ficient pressure,  the  water  molecules  re-enter  the  crystal  and  the 
original  hydrate  is  recovered.  More  remarkable  still  is  the  fact 
that,  instead  of  water  vapor,  gases  like  hydrogen  sulphide,  ammonia, 
carbon  dioxide,  and  alcohol  vapor  can  be  used,  and  enter  the  crystal 
taking  the  place  of  the  water  (see  Permutite).  These  substances 
are  chemically  so  different  from  water,  and  from  one  another,  that 
the  phenomenon  looks  more  like  a  physical  substitution  than  a 
replacement  in  chemical  combination.  These  and  other  facts  suggest 
that  hydrates  may  be  crystallographic  arrangements  of  different 
kinds  of  particles,  of  a  physical  nature,  rather  than  chemical  com- 


THE  ATOMIC  NUMBERS  469 

pounds  in  the  ordinary  sense.  This  suspicion  has  recently  been 
confirmed  in  a  remarkable  way. 

Moseley's  Atomic  Numbers.  —  We  have  seen  that  simple, 
mathematical  relations  between  the  atomic  weights  and  the  physi- 
cal or  chemical  properties  of  an  element  do  not  exist.  In  several 
instances,  the  atomic  weights  are  not  even  in  the  same  order  as 
are  the  values  of  the  properties.  We  have  now  obtained  from 
another  direction  numbers  which  seem  to  be  more  fundamental 
even  than  atomic  weights. 

Visible  light,  X-rays,  and  wireless  electric  wav  s  are  all  vibrations 
of  the  same  nature  in  the  ether.  They  differ  only  in  wave-length, 
the  order  of  the  wave-lengths  being  10~B  cm.,  10~8  cm.,  and  106  cm. 
(10  kilometers),  respectively.  Now,  just  as  the  spectrum  of  visible 
light  is  obtained  by  using  a  grat- 
ing, on  which  the  rulings  are  sep- 
arated by  distances  of  the  order 
of  the  wave-length  of  such  light, 
so  ordinary  crystals  give  spectra 
of  X-rays,  because  they  are  com-  p^  113> 

posed  of   particles   arranged   in 

rows  about  one  thousand  times  closer  and  so  form  a  suitable  grating 
for  X-rays.  This  fact  was  first  discovered  by  Dr.  Laue  of  the 
University  of  Zurich  (1912).  The  X-rays  are  produced  in  an  evac- 
uated tube  by  cathode  rays,  which  are  streams  of  electrons  emanating 
from  the  cathode  (C,  Fig.  113),  when  they  strike  the  anti-cathode 
(A). 

With  different  elements  on  the  anti-cathode,  X-rays  of  slightly 
different  wave-lengths,  and  giving  therefore  different  X-ray  spectra, 
are  produced.  By  using  different  elements,  Moseley  (1914)  has 
found  that  the  higher  the  atomic  weight  the  shorter  the  wave-length 
of  the  X-rays.  When  the  elements  are  arranged  in  ^the  order  of  these 
wave-lengths,  whole  numbers  can  be  assigned  to  each  which  are 
inversely  proportional  to  the  square  roots  of  the  wave-lengths  of 
corresponding  lines  in  their  X-ray  spectra.  These  atomic  numbers 
have  been  determined  for  most  of  the  elements,  the  atomic  weights 
of  which  lie  between  those  of  aluminium  and  uranium.  In  the  fol- 
lowing table,  the  atomic  numbers  for  these  elements  are  given  and, 
for  the  sake  of  greater  completeness,  numbers  for  the  twelve  elements 
preceding  aluminium  have  been  inserted  also. 


470 


INORGANIC  CHEMISTRY 


ATOMIC  NUMBERS   (MOSELEY) 


HI 

He     2 
Ne  10 

A     18 

Li  3 

Na  11 
K  19 
Cu  29 
Rb  37 
Ag  47 
Cs  55 
Au  79 

Gl     4 
Mgl2 
Ca  20 
Zn  30 

Sr    38 
Cd  48 
Ba  56 
Hg  80 
Ra  88 

B       5 
Al    13 

Sc    21 
Ga  31 
Y     39 
In    49 
La  57 
Tl    81 
Act  89 

C      6 

Si    14 
Ti    22 
Ge  32 
Zr   40 
Sn   50 
Ce  58 
Pb   82 
Th  90 

N      7 
P     15 
V     23 
As   33 
Cb  41 
Sb   51 
Ta73* 
Bi    83 
U-X291 

0       8 
S      16 
Cr   24 
Se    34 
Mo  42 
Te   52 
W    74 
Po   84 
U     92 

F       9 
Cl    17 
Mn25 
Br    35 
—    43 
I      53 
—    75 
—    85 

Fe   26 

Co  27 

Ni   28 

... 

Kr  36 

Ru  44 

Rh  45 

Pd  46 

... 

Xe  54 

6s   76 

Ir    77 

Pt   78 

Nt  86 

—  87 

*  The  atomic  numbers  59-72  are  those  of  the  metals  of  the  rare  earths:  Pr  59, 
Nd  60,  — 61,  Sa  62,  Eu  63,  Gd  64,  Tb  65,  Dy  66,  Ho  67,  Er  68,  Tm  69,  Yb  70, 
Lu71,  — 72. 

It  will  be  seen  that  there  is  a  whole  number  available  for  every 
known  element,  up  to  and  including  uranium,  and  not  omitting  the 
rare  elements  which  have  no  satisfactory  place  in  the  periodic  system. 
There  are  four  blank  numbers  in  the  table,  of  which  three  corre- 
spond to  spaces  below  Mn  in  the  periodic  system,  and  two  more 
amongst  the  rare  elements,  indicating  only  six  elements  with  atomic 
weights  less  than  that  of  uranium  yet  to  be  discovered.  The  atomic 
numbers  of  argon  and  potassium  place  them  in  the  chemically  cor- 
rect order,  while  the  atomic  weights  do  not.  The  same  is  true  of 
tellurium  and  iodine.  Finally,  it  is  evident  that  the  atomic  weight 
of  each  element  is,  roughly,  double  its  atomic  number. 

The  atomic  numbers  represent  the  number  of  unit  positive 
charges  of  electricity  in  the  nucleus  of  the  atom  of  each  element 
(p.  354).  Rutherford  has  shown  that  the  nucleus  contains  almost 
the  whole  mass  of  the  atom,  although  one  or  more  electrons  (nega- 
tive) are  present  also.  Thus,  the  positive  nucleus  of  the  hydrogen 
atom  is  1800  times  heavier  than  an  electron.  The  nucleus,  however, 
is  very  minute,  having  a  diameter  only  about  one-eighteen  hundredth 
of  that  of  an  electron. 

The  atomic  numbers  apparently  determine  all  the  properties 
of  each  element,  and  are  more  fundamental  than  the  atomic  weights. 
The  latter  are  secondary  properties,  in  most  cases  modified  by  other 
factors,  and  in  two  cases  actually  thrown  out  of  order  by  such  factors. 

Crystal  Structure.  —  In  this  connection  it  may  be  mentioned 
that  by  using  crystals  of  different  substances  as  X-ray  gratings, 
W.  L.  Bragg  (1914)  has  been  able  to  measure  the  distances  between 
the  rows  of  particles  in  crystals.  He  also  finds  that  the  particles, 


THE  ATOMIC  NUMBERS  471 

the  regular  arrangement  of  which  gives  the  structure  (p.  171)  of  the 
crystal  (e.g.,  a  cube  of  common  salt),  are  not  the  molecules  of  the 
compound,  much  less  aggregates  of  such  molecules,  but  the  atoms 
of  the  constituent  elements.  It  would  thus  appear  that  the  physical 
forces  (if  we  may  call  them  physical)  which  hold  the  crystalline 
solid  together  have  subordinated  the  chemical,  molecular  structure, 
and  have  arranged  the  constituent  atoms,  as  the  units  of  the 
structure,  in  a  crystallographic  pattern.  Of  course,  when  the  crystal- 
form  is  destroyed,  by  melting,  solution,  or  vaporization,  the  neigh- 
boring atoms  remain  united  in  groups,  constituting  the  chemical 
molecules  of  the  substance. 

Exercises. —  1.   Can  you  explain  the  presence  of  free  selenium 
in  the  flues  of  pyrite  burners  (p.  454)? 

2.  How  should  you  attempt  to  obtain  H^Te,  and  what  physical 
and  chemical  properties  should  you  expect  it  to  possess? 

3.  Why  does  the  existence  of  tellurium  tetrachloride  in  solution 
in  aqueous  hydrochloric  acid  show  tellurium  to  be  somewhat  metallic 
in  chemical  properties? 

4.  Make  a  list  of  bivalent  elements  and  criticize  this  method  of 
grouping  as  a  means  of  chemical  classification. 

5.  Write  down  the  symbols  of  the  elements  in  the  fourth  series 
(that  beginning  with  rubidium,  and  ending  with  iodine)  on  p.  461, 
Record  the  valence  of  each  element  toward  oxygen,  using  for  refer- 
ence the  chapters  in  which  the  oxygen  compounds  are  described. 


CHAPTER   XXIII 

OXIDES   AND   OXYGEN  ACIDS   OF   THE   HALOGENS. 
OXIDATION  AND   REDUCTION 

THE  chief  subjects  of  practical  importance  touched  upon  in  this 
chapter  are  connected  with  bleaching  powder  CaCl(OCl),  and  potas- 
sium chlorate  KC1O3  and  perchlorate  KC104.  Hence  our  attention 
will  be  largely  directed  to  the  modes  of  making  these  substances  and  to 
their  relations  to  one  another.  Incidentally,  we  shall  encounter  many 
actions  of  a  complex  and,  to  us,  more  or  less  novel  kind. 

Compounds  of  Chlorine  Containing  Oxygen.  —  The  follow- 
ing are  the  names  and  formulae  of  the  substances : 

HC1O  Hypochlorous  acid  C120  Hypochlorous  anhydride 

(HC1O2)  Chlorous  acid  

C1O2  Chlorine  dioxide 

HC1O3  Chloric  acid  

HC1O4  Perchloric  acid  C12O7  Perchloric  anhydride 

There  are  also  salts  of  these  acids,  like  the  three  substances 
mentioned  in  the  first  paragraph.  Chlorous  acid  is  itself  unknown, 
but  potassium  chlorite  KC1O2  and  some  other  derivatives  have  been 
made. 

The  two  anhydrides  (p.  150),  when  brought  into  contact  with 
water,  combine  with  it  to  form  the  acids  opposite  which  they  stand 
in  the  table.  Chlorine  dioxide  (q.v.),  however,  is  not  related  to  any 
one  acid  in  this  way. 

All  these  compounds  differ  from  most  that  we  have  hitherto  dis- 
cussed inasmuch  as  not  one  of  them  can  be  made  by  direct  union  of 
the  simple  substances. 

Nomenclature  of  the  Acids  and  their  Salts.  —  The  acids 
and  salts  are  named  on  a  plan  similar  to  that  used  in  the  case  of 
the  sulphur  acids : 

472 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS       473 

KC1O  Potassium  hypochlorite,  HC10  Hypochlorous  acid, 

KC102  Potassium  chlorite,  HC1O2  Chlorous  acid, 

KClOs  Potassium  chloral,  HClOs  Chloric  acid, 

KCICX  Potassium  perchlorate.  HC104  Perchloric  acid. 

It  should  be  noted,  however,  that  the  use  of  ic  and  ous  for  more  and 
less  oxygen,  respectively,  and  of  hypo  for  still  less  and  of  per  for 
still  more  oxygen,  are  simply  relative  terms  within  a  single  group. 
Thus,  sulphuric  acid  H2SC>4  has  a  composition  entirely  different 
from  chloric  acid,  and  both  of  these  differ  in  composition  from 
phosphoric  acid  H3PO4.  The  names  and  formulae  of  each  group 
must  be  learned,  separately. 

Chlorine   Monoxide  or   Hypochlorous  Anhydride    CIZO.  — 

A  solution  of  pure  hypochlorous  acid  is  most  easily  prepared  by 
dissolving  the  anhydride  in  water.  The  anhydride  is  obtained  by 
passing  chlorine  gas  over  warmed  mercuric  oxide*  HgO  (Fig.  88, 
p.  217) .  Each  of  the  constituents  of  the  oxide  combines  with  chlorine : 

HgO  +  2C12  ->  HgCl2  +  ClaO. 

The  mercuric  chloride  then  unites  with  another  formula-weight  of 
the  mercuric  oxide  to  form  a  solid  basic  mercuric  chloride  HgO, 
HgCl2,  which  remains  in  the  tube.  The  chlorine  monoxide  is  a 
brownish-yellow,  heavy,  easily  liquefied  gas  (b.-p.  5°).  Both  the 
gaseous  and  liquefied  forms  of  it,  the  former  when  warmed,  the 
latter  when  touched  by  paper  or  dust,  decompose  into  the  constituents 
with  explosion.  The  gas  dissolves  in  water  very  easily  (200  :  1,  by 
vol.).  The  yellow  solution  of  hypochlorous  acid  which  results: 

C12O  +  H20  *HF  2HOC1, 

has  a  strong  odor  of  chlorine  monoxide,  because  the  combination 
is  reversible,  as  it  is  in  sulphurous  acid  (p.  444).  There  are  other 
ways  of  preparing  a  dilute  solution  of  the  acid  (see  below). 

Properties  of  Hypochlorous  Acid.  —  1.  Hypochlorous  acid  is 
unstable,  and  cannot  be  made,  excepting  in  solution,  or  kept,  ex- 
cepting in  dilute  solution.  This  is  in  consequence  of  its  tendency 
to  decompose  in  three  different  ways,  one  of  which,  the  liberation 
of  the  anhydride,  has  just  been  mentioned. 

*  The  crystalline,  red  oxide  is  not  sufficiently  active.  The  oxide  must  be 
precipitated  from  sodium  hydroxide  and  mercuric  nitrate  solutions,  it  must 
be  washed  thoroughly  on  a  filter,  and  be  dried  at  300-400°  before  use. 


474  INORGANIC  CHEMISTRY 

2.  Hypochlorous  acid  is  a  little-ionized,  weak  acid. 

HOC1  +*  H+  +  C10- 

It  neutralizes  active  bases,  its  ionization  equilibrium  'being  dis- 
placed forwards  as  the  hydrogen-ion  H+  is  removed  to  form  water: 

NaOH  +  HOC1  ^±  NaOCl  +  H20. 

3.  The  solution,  if  strong,  or  when  boiled,  gives  off  chlorine 
monoxide  C12O,  the  union  with  water  being  reversible. 

4.  If  the  solution  is  concentrated,  much  of  the  hypochlorous 
acid  changes   gradually  into  chloric  acid  and  hydrogen    chloride. 
This  is  a  self -oxidation.     It  occurs  even  in  the  dark: 

3HOC1  nt  HC103  +  2HC1. 

5.  When  the  solution  is  exposed  to  sunlight,  oxygen  is  evolved. 

2HOC1  ->  2HC1  +  02  T  • 

This  decomposition  always  takes  place  in  sunlight,  whether  the 
acid  is  present  alone  in  the  water,  or  along  with  other  substances. 
We  have  already  noted  this  fact  in  discussing  chlorine-water  (p.  223), 
which  contains  this  acid. 

6.  In  consequence  of  the  ease  with  which  it  gives  up  oxygen, 
hypochlorous  acid  is  a  strong  oxidizing  agent.     In  this  direction 
i£  has  several  important  commercial  applications  (see  below). 

Commercial  Preparation  of  Hypochlorites.  —  For  commer- 
cial purposes,  pure  hypochlorites  are  not,  as  a  rule,  required. 
Hence,  sodium  or  potassium  hypochlorite  is  prepared  by  the  action 
of  sodium  or  potassium  hydroxide  on  chlorine-water.  The  latter 
contains  both  hydrochloric  and  hypochlorous  acids,  and  so  a 
solution  containing  a  mixture  of  sodium  or  potassium  chloride  and 
hypochlorite  is  obtained: 

C12  +  H2O  ^±  HC1  +  HOC1.  (1) 

HC1  +  KOH  <=>  KC1  +  H2O.  (2) 

;     HOC1  +  KOH  <=±  KOC1  +  H20.  (3) 

Although  action  (1)  is  only  partial,  being  strongly  reversible,  the 
neutralization  of  the  two  acids  in  actions  (2)  and  (3)  displaces  the 
first  equilibrium,  and  all  three  actions  proceed  to  completion. 
Action  (1),  followed  by  (2)  or  (3),  is  a  pair  of  consecutive  actions 
(p.  445),  of  which  the  second  (the  neutralization)  is  the  speedier 
of  the  two.  Both  pairs  of  consecutive  actions  (1)  +  (2)  and  (1) 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS        475 

+  (3),  can  be  combined  in  one  equation.  Thus,  omitting  the 
water,  which  appears  both  among  products  and  initial  substances 
and  in  any  case  is  present  in  large  excess  as  a  solvent,  and  omitting 
also  the  two  acids,  which  are  used  up  as  quickly  as  they  are  pro- 
duced by  equation  (1)  and  are  not  amongst  the  actual  products, 
we  get,  by  addition  of  the  three  equations  (cf.  p.  270),  the  final 
equation  : 

Cl,  +  2KOH  ->  KC1  +  KOC1  +  H20. 

As  lime  is  a  less  expensive  alkali  than  is  potassium  or  sodium 
hydroxide,  it  is  largely  used.  The  chlorine  is  led  into  rotating 
cylinders  containing  quicklime  CaO  : 


OC1 

The  product  is  not  a  mixture,  but  a  mixed  salt  (p.  401),  known  as 
bleaching  powder  or  "  chloride  of  lime."  The  fact  that  this  is  a 
mixed  salt  does  not  interfere  with  its  use  as  a  commercial  source  of 
hypochlorous  acid.  It  is  only  moderately  soluble  in  water. 

Hypochlorous  Acid  from  Bleaching  Powder.  —  1.  When 
bleaching  powder  is  dissolved  in  water,  being  a  salt,  it  is  very  ex- 
tensively ionized  (see  formulation,  below).  If  now  an  active  acid, 
that  is,  one  giving  a  large  concentration  of  hydrogen-ion,  is  added, 
the  values  of  the  products  of  the  concentrations  (H+)  X  (Cl~)  and 
(H+)  X  (OC1~),  on  which  depend  the  extent  to  which  molecules  of 
HC1  and  HOC1  will  be  formed  (p.  359),  are  large.  HC1O,  being 
little  ionized,  is  formed  extensively:  HC1,  being  highly  ionized  is 
formed  in  much  smaller  amount.  Both,  however,  interact  to  pro- 
duce chlorine  and  water,  and  this  displaces  the  other  equilibria. 
Hence  an  active  acid  decomposes  the  salt  almost  completely.  An 
active  acid  gives,  therefore,  chlorine-water,  and  not  pure  hypo- 
chlorous  acid. 

CaCl(OCl)  <=±  Ca++  +  Cl-  +  OCr       2.  A    weak    acid,    however, 

H2S04  *±  SO4=  4-  H+  +  H+     like     boric     acid     or     carbonic 

JJ          IT      acid,    gives    so    low    a    concen- 

HC1    HOC1  tration    of    H+    that    union    of 

jl  this   ion   with   OC1~   occurs   to 

T,    ^'     ~.      form    the    little    ionized    HOC1 

2     only,   and   practically   no   com- 

bination of  H+  with  Cl~  takes  place  (see  bleaching). 


476  INORGANIC  CHEMISTRY 


CaCl(OCl)  <=±  Ca++  +  Cl~  +  OCn 
H2C03  <=*  CO3=  +  H+  +  H+    r 

When  the  dilute  mixture  is  distilled,  chlorine  monoxide  (2HOC1  ^ 
H2O  -h  C120)  passes  over  with  the  steam,  and  so  a  dilute  hypo- 
chlorous  acid  can  be  obtained. 

Hypochlorous  Acid  from  Chlorine-Water.  —  An  interesting 
way  of  obtaining  dilute  hypochlorous  acid  is  to  add  chalk  CaCO3 
to  chlorine-water  and  distil.  Here,  the  chalk  is  insoluble,  and  so 
gives  a  very  low  concentration  of  Ca++  +  CO3=.  The  HC1  in  the 
chlorine-water  gives,  however,  a  sufficiently  large  concentration 
of  H+  to  combine  with  the  CO3~  to  form  H2CO3,  which  is  hardly 
ionized  at  all.  This  carbonic  acid  H2CO3  then  decomposes  and 
carbon  dioxide  is  liberated: 


TT  rn  _>  „ 
2HC1  - 


The  hypochlorous  acid,  however,  remains  molecular  HOC1,  gives 
almost  no  H+,  and  so  for  the  most  part  remains  unaffected.  It 
can  afterwards  be  distilled  off  with  the  water. 

Hypochlorous  Acid  as  an  Oxidizing  Agent.  —  Hypochlorous 
acid,  in  decomposing  into  oxygen  and  hydrochloric  acid,  gives  off 
heat.  HOC1,  Aq  —  »  HC1,  Aq  +  O  +  9300  cal.  Hence  more  energy 
is  liberated  in  oxidation  by  the  acid  than  in  oxidation  by  free 
oxygen,  and  the  acid  is  therefore  more  active  as  an  oxidizing  agent 
(p.  314).  Thus,  hypochlorous  acid,  either  in  pure  solution  or  in  the 
form  of  chlorine-water,  oxidizes  sulphurous  acid  instantly  : 

H2S03  +  HOC1  -»  H2S04  +  HC1. 

It  also  oxidizes  bromine  and  iodine,  in  water,  although  these  ele- 
ments are  not  affected  by  free  oxygen,  giving  bromic  and  iodic 
acids,  respectively: 

5HC10  +  I2  +  H20  ->  5HC1  +  2HIO3. 

The  solution  also  oxidizes  organic  colored  substances  (p.  315), 
producing  colorless,  or  less  strongly  colored  ones.  Thus,  it  oxidizes 
indigo  (deep  blue)  quickly  to  is#tin,  a  yellow  substance  relatively 
pale  in  color: 

Ci6H10N202  +  2HOC1  ->  2C8H5NO2  +  2HC1. 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS        477 

In  ways  just  as  definite  as  this,  hypochlorous  acid  will  change  the 
composition  of  other  colored  substances,  although,  since  we  do  not 
know  the  formulae  of  all  these  substances,  we  cannot  always  write 
equations  for  the  actions.  Thus,  the  interaction  by  which  chloro- 
phyll, the  green  coloring  matter  of  plants,  is  bleached  is  doubtless 
similar  to  the  above,  although  the  formulae  of  materials  concerned  are 
unknown. 

Hypochlorous  Acid  as  a  Bleaching  Agent.  —  It  is  on  account 
of  its  oxidizing  power  that  hypochlorous  acid  is  used  commercially 
in  bleaching.  It  is  not  applied  to  paints,  which  are  chiefly  mineral 
substances,  but  to  complex  compounds  of  carbon,  such  as  consti- 
tute the  coloring  matters  of  plants  and  of  those  artificial  dyes  which 
are  now  manufactured  in  great  variety.  It  should  be  understood 
that  the  great  majority  of  the  complex  compounds  of  carbon  are 
colorless.  Even  a  slight  chemical  change,  affecting  only  one  or  two 
of  the  atoms  in  a  complex  molecule,  is  thus  almost  sure  to  give  a 
colorless  or  much  less  strongly  colored  material. 

Cotton  and  linen,  in  their  original  states,  are  not  pure  white. 
Bleaching  is,  therefore,  an  extensive  and  most  important  industry. 
The  yarn  or  cloth  must  first  be  freed  from  cotton-wax  and  tannin, 
since  the  former  would  protect  it  from  the  action  of  the  bleaching 
agent,  and  both  would  make  the  subsequent  dyeing  uneven.  The 
material  is,  therefore,  first  boiled  with  very  dilute  sodium  hydroxide 
solution,  and  washed  with  water.  The  goods  are  then  saturated 
with  bleaching  powder  solution,  and  piled  loosely  until  the  coloring 
matter  has  been  oxidized.  They  are  finally  washed  with  extreme 
thoroughness. 

As  a  rule,  an  active  acid  is  not  added.  The  bleaching  is  pro- 
duced by  the  hypochlorous  acid  liberated  by  the  action  of  the  car- 
bon dioxide  from  the  air0  The  carbon  dioxide  dissolves  in  the 
water  of  the  solution  on  the  goods,  and  forms  carbonic  acid: 
C02  +  H2O^H2C03  (see  p.  476,  par.  1).  The  subsequent  wash- 
ing removes  all  traces  of  the  bleaching  powder,  of  the  lime  which 
the  powder  often  contains,  and  of  the  hypochlorous  acid,  which 
otherwise  would  act  gradually  upon  the  cotton  or  linen  and  "rot" 
it.  Bleaching  agents,  when  used  in  the  household  without  sufficiently 
careful,  subsequent  washing,  are  liable  to  cause  serious  damage 
from  this  cause. 

Cotton  and  linen  are  composed  of  cellulose  (CeHuA)*,  a  rather 
inert  substance,  and  one  which  is  very  slowly  acted  upon  by  dilute 


478  INORGANIC  CHEMISTRY 

hypochlorous  acid.  Hence,  with  brief  contact  and  proper  han- 
dling, no  damage  is  done.  Wool,  silk,  and  feathers,  however,  are 
composed  largely  of  compounds  (proteins)  containing  nitrogen 
(up  to  15  per  cent),  in  addition  to  the  above  three  elements.  Their 
constituent  material  interacts  as  easily  with  hypochlorous  acid  as 
do  the  traces  of  coloring  substances.  Hence,  since  the  fabric  itself 
would  be  attacked  by  this  agent,  sulphur  dioxide  or  sulphurous 
acid  (p.  445)  is  used  for  bleaching  these  materials. 

It  should  be  understood  that  a  cold  dilute  solution  of  hypochlorous 
acid  may  be  kept  (in  the  dark)  almost  indefinitely  and  will  not  give 
up  its  oxygen  spontaneously.  The  transfer  takes  place  when,  and 
only  when,  the  acid  comes  in  contact  with  some  substance  capable 
of  uniting  with  oxygen. 

Bleaching  Powder  in  Sanitation.  —  A  disinfectant  is  a  sub- 
stance which  destroys  bacteria  and  other  minute  organisms. 
Bleaching  powder  has  a  distinct  odor  of  chlorine  monoxide  (not 
chlorine).  This  is  due  to  the  action  of  atmospheric  carbon  di- 
oxide liberating  hypochlorous  acid  (p.  476).  The  dry  powder 
therefore  will  disinfect  the  air  and  surrounding  objects.  It  must 
be  used  with  discretion,  however,  as  the  gas  is  very  corrosive. 

As  already  mentioned  (p.  142),  in  the  purification  of  city  waters 
the  organisms  (associated  in  some  way  with  colon  bacilli)  which  give 
rise  to  typhoid  fever  are  destroyed  by  adding  a  small  proportion  of 
bleaching  powder  in  the  form  of  a  2  per  cent  solution  (about  17-24 
Ibs.  of  the  powder  per  million  gallons  of  water).  The  salt  is  hydro- 
lyzed  (p.  398),  giving  a  basic  calcium  chloride  and  free  hypochlorous 
acid.  The  latter  kills  the  organisms,  and  is  itself  decomposed  in 
the  process,  so  that  nothing  offensive  remains  in  the  water.  There 
is  only  a  minute  increase  in  the  proportion  of  salts  of  calcium  (hard- 
ness). Sewage  is  sometimes  freed  from  pathogenic  organisms  in 
the  same  way. 

Recently,  chlorine-water,  made  by  use  of  cylinders  of  liquid 
chlorine  (p.  221),  has  in  many  cases  taken  the  place  of  bleaching 
powder  solution  for  this  purpose. 

Chlorine  not  a  Bleaching  Agent.  —  Chlorine  itself  is  often, 
erroneously,  described  as  a  bleaching  agent.  If  a .  dry,  colored 
cloth  be  hung  for  a  week  in  chlorine  gas,  dried  by  a  little  sulphuric 
acid  in  the  bottom  of  the  bottle  (Fig.  114),  little  or  no  change  in  the 
color  will  occur.  But  a  wet  rag  is  bleached  as  soon  as  the  chlorine 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS        479 

has  time  to  dissolve  in  the  water  and  give  the  necessary  hypochlorous 
acid.     Flowers  are  bleached  by  dry  chlorine  gas,  be- 
cause by  their  nature  they  contain  the  indispensable 
water. 

Thermochemistry  of  Hypochlorous  Acid.  —  As 

we  have  seen  (p.  35),  chemical  changes  which  proceed 

spontaneously  are  accompanied  by  a  transformation 

of  free  internal  energy  into  some  other  form  of  energy. 

Hence,  a  substance,  or  system  of  substances,  which 

undergoes    such  a  change,    possesses   more    chemical 

energy  and  activity  before  the  change  than  after  it. 

In  consequence,  if  some  given  chemical  change  uses  the    _ 

products  of  such  an  action,  and  can  be  brought  about        FlQ<  114 

by  the  employment  of  the  original  substance,  the  em- 

ployment of  the  latter  will  involve  a  greater  liberation  of  energy, 

and  will  therefore  be  more  likely  to  secure  the  consummation  of  the 

change  in  question. 

The  decomposition  of  hypochlorous  acid  and  of  chlorine  monoxide 
are  cases  where  there  is  a  very  marked  difference  between  the  amount 
of  chemical  energy  in  the  original  substances  and  in  the  products  of 
decomposition,  hydrogen  chloride  and  free  oxygen  in  the  first  case,  and 
free  chlorine  and  oxygen  in  the  second.  Hence  the  changes  into  these 
substances  sometimes  are  of  the  nature  described  as  explosive.  A 
more  important  fact,  however,  is  that,  on  this  account,  hypochlorous 
acid  and  chlorine  monoxide  are  more  active  oxidizing  agents  than  is 
free  oxygen  gas.  The  energy  liberated  in  the  decomposition  of  the 
hypochlorous  acid  has  to  be  added  (p.  315)  to  that  which  free  oxygen 
could  give,  if  performing  the  same  oxidation,  in  order  that  the  total 
fall  in  energy,  which  measures  the  tendency  of  the  action  to  take 
place,  may  be  estimated.  Hence,  substances  that  are  not  affected  by 
free  oxygen  may  be  changed  instantly  by  hypochlorous  acid.  This 
explains,  for  example,  the  oxidation  by  hypochlorous  acid  of  many 
carbon  compounds,  including  those  which  are  colored,  when  atmos- 
pheric air  is  without  action.  Thus,  the  heat  liberated  in  the  oxidation 
of  indigo  to  isatin  by  oxygen  gas,  if  it  could  be  carried  out,  would  be 
1800  cal.  The  much  greater  heat  liberated  when  hypochlorous  acid 
is  used,  we  obtain  by  adding  the  thermochemical  equations: 

2HC1O  =  2HC1  +  2O  +  18,600  cal. 
2O  =  2C8H6N02     +  1800  cal. 


_ 
C16H10N202  +  2HC1O  =  2C8H5N02  +  2HC1  +  20.400  cal. 


480 


INORGANIC  CHEMISTRY 


The  following  thermochemical  equations  give  a  rough  idea  of  the  relative 
oxidizing  powers  of  the  chief  oxygen  acids  of  the  halogens: 


HC1O,  Aq  =  HC1,  Aq  +  O  +  9,300  cal.,  or  +  9,300 
HC1O3,  Aq  =  HC1,  Aq  +  3O  +  15,300  cal.,  or  +  5,100 
HC1O4,  Aq  =  HC1,  Aq  +  4O  +  700  cal.,  or  +  170 
HBrO3,  Aq  =  HBr,  Aq  +  3O  +  15,000  cal.,  or  +  5,000 
HIO3,  Aq  =  HI,  Aq  +  3O  —  42,900  cal.,  or  —  14,300 
HIO4,  Aq  =  HI,  Aq  + 40-34,500 cal.,  or-  8,600 


cal.  for 
each  atomic 
weight  of 
oxygen. 


Formerly  a  different  explanation  for  actions  like  that  of  hypochlorous  acid, 
when  it  behaves  as  an  oxidizing  agent,  was  offered.  It  was  suggested  that  the 
oxygen  was  first  liberated  from  the  acid:  HOC1  — »  HC1  +  O,  and  that  the  single 
atoms  of  the  element  so  produced  were  more  active  than  molecular  oxygen. 
Exactly  how  the  hypochlorous  acid,  lacking  all  consciousness,  knew  that  an 
oxidizable  body  was  present,  and  proceeded  to  liberate  the  atoms  of  oxygen,  was 
never  explained.  This  oxygen,  which  was  supposed  to  interact  in  the  moment  of 
its  production,  was  called  nascent  oxygen.  But  it  will  be  seen  that  such  an  ex- 
planation is  entirely  unnecessary.  The  activity  of  the  hypochlorous  acid  on 
account  of  its  large  store  of  free  energy  sufficiently  accounts  for  the  facts  (see 
Nascent  hydrogen). 

Chemical  Properties  of  Hypochlorites.  —  When  hypochlo 
rites  are  heated  they  change  into  chlorates  (see  below).  They  may 
also  give  off  oxygen,  2CaCl(OCl)  -»  2CaCl2  +  O2.  Although  this 
decomposition  is  slow  in  cold  solutions  of  hypochlorites,  or  when 
they  are  preserved  in  the  dry  form,  it  may  be  hastened  by  means 
of  catalytic  agents.  The  addition  of  a  little  cobalt  hydroxide  (q.v.) 
to  a  paste  of  bleaching  powder  and  water  causes  rapid  evolution  of 
oxygen. 

Chlorates.  —  Like  hypochlorous  acid  itself,  the  hypochlorites 
turn  into  chlorates.  Thus,  when  chlorine  is  passed  into  a  warm, 
concentrated  solution  of  potassium  hydroxide,  and  particularly 
when  an  excess  of  chlorine  is  used,  the  potassium  hypochlorite 
changes  into  potassium  chlorate  KClOa  as  fast  as  it  is  formed : 


3KC10  ->  KC1O3  +  2KC1. 


(1) 


To  secure  the  three  molecules  of  the  hypochlorite,  the  equation  for- 
merly given  (p.  475)  must  be  tripled: 


3C12  +  6KOH  ->  3KC1  +  3KC1O  +  3H2O. 


(2) 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS        481 

When  these  are  added,  and  the  intermediate  substance  is  left  out,  the 
final  equation  is  obtained: 

3C12  +  6KOH  -»  KC103  +  5KC1  +  3H20. 

When  the  solution  is  cooled,  the  less  soluble  chlorate  crystallizes  out. 

This  action  involves  converting  five-sixths  of  the  valuable  potas- 
sium hydroxide  into  the  relatively  less  valuable  potassium  chloride. 
Hence,  in  practice,  the  makers  carry  out  the  corresponding  action 
with  calcium  hydroxide.  They  then  add  potassium  chloride  to  the 
resulting  solution,  containing  calcium  chloride  and  calcium  chlorate 
Ca(ClO3)2  (sol'ty  175:100  Aq,  18°).  The  potassium  chlorate 
(soPty  6.6  :  100  Aq,  18°),  formed  by  double  decomposition,  crystallizes 
when  the  solution  is  cooled. 

All  chlorates  are  at  least  moderately  soluble  in  water  (see  Table 
inside  front  cover).  Potassium  chlorate  is  used  in  making  fireworks, 
explosives,  and  matches.  An  intimate  mixture  with  sugar  C^H^Ou 
burns  with  semi-explosive  violence,  the  oxygen  of  the  salt  combining 
with  the  carbon  and  hydrogen  to  form  carbon  dioxide  and  water. 

The  Separation  of  Substances  by  their  Solubility. — When 
neither  of  the  produces  of  an  action  approaches  absolute  insolubility,  a  separation 
may  nevertheless  be  effected  more  or  less  perfectly  by  taking  advantage  of  differ- 
ence in  solubility.  Thus,  in  the  practical  method  of  making  potassium  chlorate, 
the  calcium  chloride  is  exceedingly  soluble,  while  the  potassium  chlorate  is  only 
moderately  so.  Then,  too,  the  solubility  of  the  latter  decreases  rapidly  as  the 
temperature  is  lowered  (Fig.  79,  p.  191).  Hence,  it  is  found  that  when  the  mix- 
ture is  cooled  to  — 18°  only  about  13.5  g.  of  potassium  chlorate  remain  dissolved 
in  each  liter,  and  are  lost.  At  0°  the  loss  would  be  greater,  for  at  this  temperature 
a  liter  of  pure  water  would  hold  33.3  g.,  and  a  liter  of  this  solution  would  contain 
more  than  this  on  account  of  the  uncompleted  reversible  action  (cf.  p.  381) : 

Ca(C103)2  +  2KC1  *±  CaCl2  +  2KC1O3. 

It  will  be  seen  that  we  reason  as  if  the  solubility  of  each  substance  was  inde- 
pendent of  the  presence  of  other  dissolved  bodies  (p.  187). 

By  the  use  of  this  principle,  and  the  data  in  regard  to  solubility  in  Fig.  79 
(p.  191),  a  rough  idea  may  be  obtained  of  what  may  be  expected  in  any  given 
case.  From  the  diagram  the  solubilities  at  any  given  temperature  may  be  read. 
Suppose,  for  example,  the  question  is  in  regard  to  the  quantity  of  potassium 
chlorate  we  may  expect  to  obtain  from  3  g.  of  potassium  hydroxide  dissolved  in 
7  g.  of  water  (a  30  per  cent  solution).  From  the  equation: 

3C12  +  6KOH  ->  KC1O3  +  5KC1  +  3H2O 

6X66  122.5        5X74.5 


482 


INORGANIC   CHEMISTRY 


we  find  that  336  g.  of  potassium  hydroxide  give  122.5  g.  of  chlorate  and  372.5  g. 
of  chloride.  Hence,  by  proportion,  3  g.  will  give  about  1  g.  and  3  g.  respectively. 
The  solubility,  read  from  the  diagram,  is  the  amount  of  the  salt  dissolved  by  100 
c.c.  of  water,  for  example,  56.5  g.  of  potassium  chloride  at  100°.  Some  of  the 
results  are  given  in  the  form  of  a  table: 


POTASSIUM 
CHLORIDE. 

POTASSIUM 
CHLORATE. 

Amount  formed  from  3  g  KOH       .       ... 

3  0 

1  0 

Solubility  at  )  100  c.c.  Aq  
100°  in         J      7  c.c.  Aq  
Solubility  at  )  100  c.c.  Aq  
20°  in           J      7  c.c.  Aq  

56.5 
4.0 
34.7 
2.5 

56.5 
4.0 
7.5 
0.5 

Solubility  at  f  100  c.c.  Aq  

28.0 

3.3 

0°  in             f     7  c.c.  Aq  

2.0 

0.25 

Thus,  at  20°,  at  least  2.5  g.  of  the  3  g.  of  potassium  chloride  will  remain  dissolved, 
while  half  of  the  potassium  chlorate  will  crystallize  out.  If  the  solubilities  are 
examined,  it  will  be  seen  that  the  potassium  chlorate  is  even  more  easily  obtain- 
able in  pure  condition  when  calcium  chloride  takes  the  place  of  potassium  chloride. 

Chloric  Acid.  —  Since  none  of  the  acids  of  this  series  can  be 
obtained  by  direct  union  of  their  elements  (p.  472),  it  is  usual  first 
to  prepare  the  salts,  and  to  make  the  acids  from  the  salts  by  double 
decomposition.  This  acid  may  be  obtained  in  solution  in  water, 
by  adding  the  calculated  amount  of  hydrofluosilicic  acid  to  a  solution 
of  potassium  chlorate: 


2KC103  +  H2SiF6  <=±  K2SiF6 


2HC103. 


The  potassium  fluosilicate,  being  insoluble,  is  removed  by  filtration. 
It  will  be  noted  that  double  decomposition,  involving  precipitation, 
may  thus  be  used  for  obtaining  a  soluble  product,  as  well  as  an  in- 
soluble one  (cf.  selenic  acid,  p.  454). 

The  statement  commonly  made  that  chloric  acid  is  prepared  by 
adding  dilute  sulphuric  acid  to  barium  chlorate  solution:  Ba(C103)2 
+  H2S04  ^  BaSO4  1  +  2HC103,  is  interesting.  Barium  chlorate 
is  itself  made  from  chloric  acid  and  barium  hydroxide  !  The  action  of 
chlorine  upon  a  solution  of  the  latter  substance  cannot  be  used, 
because  the  chlorate  and  chloride  of  barium  are  equally  soluble 
(see  Table),  and  cannot  be  separated  by  partial  crystallization. 

The  solution  of  chloric  acid  may  be  concentrated  (to  about  40 
per  cent)  by  evaporation,  but  must  not  be  heated  above  40°,  as  the 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS        483 

acid  decomposes  near  this  temperature.  The  resulting  thick,  color- 
less liquid  has  powerful  oxidizing  qualities,  setting  fire  to  paper 
(made  of  cellulose  CeHioCX)  which  has  been  dipped  into  it.  It  con- 
verts iodine  into  iodic  acid,  5HC103  +  3I2  +  3H2O  ->  6HIO3  + 
5HC1.  When  warmed  beyond  40°  the  acid  decomposes,  giving 
chlorine  dioxide  and  perchloric  acid: 

3HC1O3  ->  H20  +  2C102  +  HC1O4  (see  p.  495). 

Chlorine  Dioxide:  Chlorous  Acid.  —  Chlorine  dioxide  C102 
(see  above)  is  a  yellow  gas  which  may  be  liquefied,  and  boils  at 
+  10°.  The  gas  and  liquid  are  violently  explosive,  the  substance 
being  resolved  into  its  elements  with  liberation  of  much  heat.  It  is 
formed  whenever  chloric  acid  is  set  free,  and  hence  it  is  seen  when  a 
little  powdered  potassium  chlorate  is  touched  with  a  drop  of  con- 
centrated sulphuric  acid  (end  of  last  section).*  Concentrated 
hydrochloric  acid  turns  yellow  from  the  same  cause  when  any  chlorate 
is  added  to  it.  These  actions  are  used  as  tests  for  chlorates,  and 
distinguish  them  from  perchlorates  (q.v.).  With  water,  chlorine 
dioxide  gives  a  mixture  of  chlorous  acid  HC102  and  chloric  acid,  and 
with  bases  a  mixture  of  the  chlorite  and  chlorate. 

Perchlorates.  —  When  heated,  chloric  acid  and  chlorates  give 
perchloric  acid  (p.  483)  and  perchlorates  respectively.  The  chlorates 
also  give  oxygen  at  the  same  time  (p.  83) : 

( 2KC1O3  -»  2KC1  +  302, 
1 4KC103  -»  3KC1O4  +  KC1. 

These  actions,  like  the  three  decompositions  of  hypochlorous 
acid  (p.  474),  are  independent,  and  proceed  simultaneously.  They 
are  concurrent  reactions  (p.  485).  Their  relative  speed,  however, 
varies  with  the  temperature,  and  the  decomposition  into  chloride 
and  oxygen  may  completely  outrun  the  other  when  a  catalytic 
agent  like  manganese  dioxide,  which  hastens  only  one  of  the  two 
actions,  is  added  (p.  97).  When  pure  potassium  chlorate  is  heated 
cautiously,  about  one-fifth  of  it  has  lost  all  its  oxygen  by  the  time 
the  rest  has  turned  into  perchlorate.  The  mixture  may  be  separated 

*  The  mixture  of  sugar  and  potassium  chlorate  (p.  481)  can  be  set  on  fire  by  a 
drop  of  sulphuric  acid  [Lect.  exp.].  The  latter  liberates  chloric  acid,  which  in 
turn  gives  C1O2,  and  the  latter,  being  a  violent  oxidizing  agent,  starts  the  com- 
bustion of  the  sugar. 


484  INORGANIC   CHEMISTRY 

by  grinding  with  the  minimum  quantity  of  water  which  will  dissolve 
the  chloride  it  contains.  The  perchlorate,  having  at  15°  less  than 
one-twentieth  of  the  solubility  of  the  chloride,  will  remain,  for  the 
most  part,  undissolved.  The  perchlorates  are  much  more  stable 
(p.  148)  than  the  chlorates,  or  hypochlorites  :  they  are  all  soluble 
in  water,  and  they  are  used  in  making  matches  and  fireworks. 


Perchloric  Acid  HCIO^  and  Perchloric  Anhydride 

Pure  perchloric  acid  explodes  when  heated  above  92°.  But,  like 
other  liquids,  its  boiling-point  is  lower  when  its  vapor  is  under  reduced 
pressure  (cf.  p.  146).  At  56  mm.  pressure  it  boils  at  39°,  a  temperature 
at  which  hardly  any  decomposition  is  noticeable.  Hence  the  acid 
may  be  made  by  mixing  potassium  perchlorate  and  concentrated 
sulphuric  acid  and  distilling  the  mixture  cautiously  in  a  vacuum 
(p.  316). 

KC104  +  H2S04  *±  KHS04  +  HC1O4T. 

Perchloric  acid  is  a  colorless  liquid,  which  decomposes,  and  often 
explodes  spontaneously,  when  kept.  A  70  per  cent  solution  in  water 
is  perfectly  stable,  however.  Although  it  is  an  active  oxidizing 
agent,  it  is  not  so  active  as  chloric  acid,  and  does  not  oxidize  hydrogen 
chloride  in  cold  aqueous  solution.  Hence  a  drop  of  hydrochloric 
acid  placed  on  a  crystal  of  a  perchlorate  gives  no  yellow  color.  When 
the  acid  is  liberated  by  concentrated  sulphuric  acid,  it  does  not  at 
once  give  the  yellow  chlorine  dioxide  (p.  483)  . 

Perchloric  anhydride  C1207  may  be  prepared  by  adding  phosphoric 
anhydride  to  perchloric  acid  in  a  vessel  immersed  in  a  freezing  mix- 
ture, P205  +  2HC1O4  ->  2HP03  +  C1207.  Phosphoric  anhydride  is 
often  used  in  this  way  for  removing  the  elements  of  water  from 
compounds.  It  combines  with  water  to  form  metaphosphoric  acid 
HPO3.  By  gently  warming  the  mixture,  the  perchloric  anhydride 
can  be  distilled  off.  It  is  a  colorless  liquid  boiling  at  82°  (760  mm.), 
and  exploding  when  struck  or  too  strongly  heated. 

Relation  of  Anhydride  and  Acid  or  Salt.  —  The  derivation 
of  the  formula  of  the  anhydride  from  Lthat  of  the  acid  or  salt 
should  receive  special  attention.  In  the  mind  of  the  chemist,  the 
one  always  instantly  suggests  the  other,  so  often  does  he  think 
of  them  as  potentially  the  same  substance.  The  beginner,  how- 
ever, finds  this  habit  hard  to  acquire,  and  indeed  is  more  likely  to 
blunder,  in  trying  to  divide  the  formula  of  an  acid  into  the  formulae 
of  water  and  the  anhydride,  than  in  any  other  calculation  he  makes. 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS        485 

The  rule  is:  If  the  formula  of  the  acid  shows  an  even  number 
of  hydrogen  atoms  (H2S04  or  H^iOJ,  subtract  all  the  elements 
of  water  (H2O  or  2H2O),  and  the  balance  is  the  anhydride  (SO3 
or  SiO2).  The  divided  formulae  are  H20,S03  or  2H2O,SiO2.  If 
there  is  an  odd  number  of  hydrogen  atoms  (HC1O4  or  HsPO^ 
double  the  formula  (H2Cl2Os  or  H6P208),  and  subtract  all  the  ele- 
ments of  water  as  before  (C1207  or  P205).  Then  check  the  result, 
by  adding  the  water  again,  and  dividing  by  two,  correcting  the 
blunder  if  one  has  been  made. 

If  the  substance  is  a  salt  (CuSO4  or  KC104),  subtract  the  oxide 
of  the  metal  (CuO  or  K2O),  taking  care  to  assign  to  the  metal  the 
same  valence  in  the  oxide  as  it  shows  in  the  salt  (SO3  or  C12O7). 

There  are  several  uses  for  this  art  of  ascertaining  the  anhydride 
corresponding  to  a  given  salt  or  acid.  One  is  in  the  making  of 
equations  (e.g.,  pp.  496,  535).  Another  is- in  finding  the  valence  of 
the  non-metal.  Thus;  in  KC1O4  the  anhydride  is  C12O7,  and  the 
valence  of  the  chlorine  is  seven.  In  H3PO4  the  anhydride  is  P2Os 
and  the  phosphorus  quinquivalent.  In  HPO3  (metaphosphoric 
acid),  the  anhydride  is  again  P20s,  and  the  phosphorus  is  therefore 
in  the  same  state  of  oxidation  —  both  are  phosphoric  acids. 

Simultaneous  Chemical  Changes  in  the  Same  Substances. 

—  When  two  or  more  reactions  go  on  simultaneously  in  the  same 
materials,  the  actions  may  be  consecutive  (p.  445)  or  they  may  be 
parallel.  In  the  latter  case  they  are  called  concurrent  reactions. 
Thus,  hypochlorous  acid  undergoes  three  different  changes: 

2HC10  ->  H2O  +  C120. 
3HC1O  ->  HC103  +  2HC1. 
2HC1O  -»  2HC1  +  O2. 

Some  molecules  decompose  into  water  and  chlorine  monoxide  (p. 
474),  while  others  give  chloric  acid  and  hydrogen  chloride,  and  still 
others  hydrogen  chloride  and  oxygen.  Since  the  same  molecule 
cannot  undergo  more  than  one  of  these  different  changes,  it  follows 
that  the  actions  are  independent  of  one  another.  This  is  shown 
by  the  fact  that  in  sunlight  the  third  predominates,  while  in  the 
dark  it  falls  far  behind  the  second.  Since  the  relative  quantities 
of  the  products  vary,  the  several  simultaneous  actions  cannot  be  put 
in  the  same  equation.  The  fundamental  property  of  an  equation  is 
to  show  the  constant  proportions  by  weight  between  every  pair  of 
substances  in  it.  Hence  three  separate  equations  are  required  in 


486  INORGANIC  CHEMISTRY 

the  present,  and  in  all  similar  cases  where  all  the  proportions  are 
not  constant.  Thus,  again,  in  the  decomposition  of  potassium 
chlorate  by  heating  (p.  483),  it  would  be  misleading  and  wrong  to 
add  the  two  equations  together  and  write,  for  the  whole  action: 

2KC1O3  ->  KC1  +  KC104  +  O2. 

This  equation  would  mean  that  the  proportions  amongst  the  prod- 
ucts were  always  KC1  :  KC1O4  :  O2  or  74.6  : 138.6  :  32,  whereas, 
in  fact,  the  proportions  vary  with  the  conditions  —  the  tempera- 
ture used  or  the  presence  of  a  catalyst  which  hastens  one  action 
but  not  the  other. 

Consecutive  reactions  (p.  445),  however,  like  (1)  followed  by  (2) 
on  pp.  433,  474,  may  be  combined  in  one  equation,  since  in  them 
all  the  proportions  must  necessarily  be  constant.  These  equations 
are  interlocked,  for  (2)  consumes  what  (1)  produces. 

Oxygen  Acids  of  Bromine.  —  No  oxides  of  bromine  have  been 
made,  but  the  acids  HBrO  (hypobromous  acid)  and  HBrO3  (bromic 
acid)  and  their  salts  are  familiar. 

By  the  action  of  bromine  on  dilute,  cold  potassium  hydroxide 
solution,  the  bromide  and  hypobromite  are  formed: 

Br2  +  2KOH  ->  KBr  +  KBrO  +  H2O. 

When  the  solution  is  heated,  the  hypobromite  turns  into  potassium 
bromate  KBr03  and  bromide.  The  actions  are  exact  parallels  of 
the  corresponding  ones  for  chlorine  (pp.  474,  480). 

Aqueous  bromic  acid  HBr03  may  be  made  in  the  same  way  as 
chloric  acid  (p.  482),  or  by  the  action  of  chlorine- water  on  bromine: 

5HC10  +  Br2  +  H20  ->  2HBr03  +  5HC1. 

The  solution  is  colorless  and  has  powerful  oxidizing  properties.  Thus, 
it  converts  iodine  into  iodic  acid:  2HBrO3  +  I2  — > 2HIO3  +  Br2. 
It  appears,  therefore,  that  iodine  has  more  affinity  for  oxygen  than 
has  bromine. 

Oxide  and  Oxygen  Acids  of  Iodine.  —  The  following  are  the 
acids  and  their  corresponding  salts: 

[HIO      Hypoiodous  acid],  [KIO  Potassium  hypoiodite], 

HIO3      Iodic  acid,  KIO3  Potassium  iodate, 

[HIO4     Periodic  acid],  NaIO4  Sodium  periodate, 

Periodic  acid,  Na2H3IOe  Disodiura  periodate. 


OXIDES  AND  OXYGEN  ACIDS  OF  THE  HALOGENS       487 

The  substances  in  parenthesis  have  not  been  isolated.     There  is  one 
oxide,  iodic  anhydride  I2Os. 

lodates  and  Iodic  Acid.  —  The  potassium  and  sodium  salts 
of  iodic  acid  are  found  in  Chile  saltpeter.  They  may  be  made,  in 
much  the  same  fashion  as  are  the  chlorates  and  bromates  (p.  481), 
by  adding  powdered  iodine  to  a  hot  solution  of  potassium  or  sodium 
hydroxide.  There  is  evidence  that  hypoiodites  are  formed  in  cold 
solutions,  but  they  change  quickly  to  iodates.  It  is  disodium  perio- 
date  Na^HsIOe,  however,  which,  being  the  least  soluble,  crystallizes 
out. 

Iodic  acid  HI03  is  formed  by  passing  chlorine  through  powdered 
iodine  suspended  in  water.  The  action  is  parallel  to  that  of  chlorine 
on  bromine  water.  A  still  better  way  is  to  boil  iodine  with  aqueous 
nitric  acid  (q.v.).  The  latter  gives  up  oxygen  readily,  and  is  here 
used  solely  on  this  account.  Hence  it  may  be  omitted  from  the 
equation,  only  the  oxygen,  of  which  it  is  the  source,  appearing: 

I2  +  H20  +  50  ->  2HI03. 

In  both  these  actions  the  initial  substances  (including  the  excess  of 
nitric  acid)  and  the  products,  with  the  exception  of  the  iodic  acid 
itself,  are  all  volatile.  When  the  solution  is  concentrated  by  evapora- 
tion, the  iodic  acid  crystallizes.  It  is  a  white  solid,  perfectly  stable  at 
ordinary  temperatures,  and  can  be  kept  indefinitely.  At  170°  it 
begins  to  give  off  water  vapor  2HI03  <=^  H20  +  I205,  leaving  iodic 
anhydride.  The  latter  is  a  white  crystalline  powder  which  may  be 
raised  to  300°  before  it,  in  turn,  breaks  up,  giving  iodine  and  oxygen. 
In  aqueous  solution  iodic  acid  is  an  oxidizing  agent,  but  does  not 
part  with  its  oxygen  so  readily  as  do  chloric  acid  and  bromic  acid.  It 
oxidizes  hydrogen  iodide  in  dilute  solution:  HIO3  +  5HI  — »  3H2O  -f- 
3I2,  all  the  iodine  being  liberated.  In  this  respect  it  resembles  concen- 
trated sulphuric  acid  (p.  438).  Dilute  sulphuric  acid  shows  no 
oxidizing  qualities. 

Various  Acids  Derived  front  One  Anhydride.  —  Some  acids 
are  related  to  their  anhydrides  as  are  hypochlorous  acid  (p.  473)  and 
sulphurous  acid  (p.  444).  One  molecule  of  the  anhydride  combines 
with  one  molecule  of  water.  In  other  cases,  however,  the  proportion 
of  water  may  be  less  or  greater  than  this.  Now  if  periodic  acid  were 
of  the  former  type  (H2O,I2O7  =  2HIO4),  its  formula  would  be 
HI04.  It  does  form  salts  of  this  type,  such  as  NalC^  and  AglO*. 


488  INORGANIC  CHEMISTRY 

But  the  free  acid  is  a  deliquescent  solid  of  the  formula 
(=  5H20,I2O7),  and  the  most  easily  prepared  salt  belongs  to  this 
type.  All  types  are  called  periodates,  however,  because  their  compositions 
are  all  founded  upon  the  same  anhydride.  The  latter  has  not  itself 
been  made.  We  usually  speak  of  various  acids  and  salts  as  being 
derived  from  the  same  anhydride,  the  word  "  derived"  being  used 
in  a  figurative  and  not  a  literal  sense. 

The  difference  between  two  acids  HIO4  and  H5IO6  is  not  at  all  the 
same  as  between  HIO3  and  HI04.  The  latter  would  represent  differ- 
ent stages  of  oxidation,  being  derived  from  I205  and  I207,  respectively, 
and  accordingly  would  be  named  iodic  acid  and  periodic  acid.  The 
former  differ  only  by  2H2O,  and  an  addition  or  subtraction  of  the 
elements  of  water  in  equivalent  quantities  is  neither  oxidation  nor 
reduction.  Hence  they  are  both  periodic  acids  (see  Phosphoric  acid). 

Periodates  and  Periodic  Acid.  —  Sodium  periodate  NaIO4  is 
found  in  Chile  saltpeter.  When  sodium  iodate  NaIO3  is  dissolved 
along  with  sodium  hydroxide  in  water,  and  chlorine  is  passed  into  the 
mixture,  the  sodium  hypochlorite  formed  from  the  latter  oxidizes  the 
iodate  NaIO3  +  O  —  >  NaIO4.  But  the  somewhat  insoluble  salt  which 
crystallizes  out  is  Na^HsIOe: 

NaI03  +  0  +  NaOH  +  H20 


Other  salts  may  be  made  from  this  one. 

An  aqueous  solution  of  periodic  acid  is  obtained  by  the  action  of 
sulphuric  acid  on  barium  periodate,  followed  by  filtration.  A 
white,  very  soluble  solid  HsIOe  remains  when  the  liquid  is  evaporated. 
When  this  is  heated,  water  and  oxygen  are  both  given  off,  and  iodine 
pentoxide  I205  alone  remains. 

Chemical  Relations.  —  The  compounds  of  the  halogens  with 
metals  and  with  hydrogen  diminish  in  stability,  with  ascending  atomic 
weight  of  the  halogen,  in  the  order:  F  (19),  Cl  (35.5),  Br  (80),  I  (127). 
Each  halogen  will  displace  those  following  it  from  this  kind  of  combi- 
nation. In  the  case  of  the  oxygen  compounds,  the  order  of  stability 
is  just  the  reverse,  those  of  iodine,  for  example,  being  the  only  ones 
which  are  reasonably  stable. 

Amongst  the  oxygen  acids  of  any  one  halogen,  those  containing 
most  oxygen  are  most  stable.  The  salts  are  in  all  cases  more  stable 
by  far  than  the  corresponding  acids. 

The  halogens  when  combined  with  metals  or  hydrogen  are  univa- 


OXIDES  AND  OXYGEN   ACIDS  OF  THE  HALOGENS        489 

lent  (HI,  KC1,  etc.).  It  is  clear,  however,  that,  when  united  with 
oxygen,  their  valence  is  higher.  The  maximum  is  shown  in  perchloric 
anhydride  C^Oy,  where  chlorine  is  septivalent. 

The  formulae  of  the  acids  might  be  written  so  as  to  retain  the 
uni  valence : 

H  -  Cl,  •  H  -  0  -  Cl,     H  -  O  -  O  -  Cl,     H-0-0-0-C1, 
H-O-O-O-O-C1. 

But  compounds  in  which  we  are  compelled  to  believe  that  two  oxygen 
units  are  united  are  usually  unstable  (e.g.,  hydrogen  peroxide,  H  —  O 
—  O  —  H),  and  we  should  expect  the  instability  would  be  greater 
with  three  and  with  four  units  of  oxygen  in  combination.  Here, 
however,  the  reverse  state  of  affairs  must  be  taken  account  of  in  our 
formulae,  for  HC1O4  is  the  most  stable  of  the  chlorine  set.  This 
reasoning,  together  with  the  septivalence  in  C12O7,  leads  us  to  assume 
the  valence  seven  in  perchloric  acid  (see  Periodic  system).  The 
structural  formulae  (cf.  p.  442)  of  some  of  these  substances  are  there- 
fore often  written  as  follows: 

O  O 

II  II 

H-C1,        H-O-C1,        H-O-C1  =  O,        Na-0-I  =  0. 

II  II 

0  O 

The  Specification  of  Chemical  Properties.  —  The  chemical  prop- 
erties of  a  substance  are  frequently  specified  so  loosely  that  the  student  has  no 
definite  guide  before  him.  The  following  illustrates  the  scheme  that  should  be 
kept  in  mind. 

1.  Stability  (p.  148),  particularly  if  a  compound,  but  applicable  also  to 
elements  like  iodine  (p.  276).     Mention  products  of  decomposition. 

2.  Molecular  weight,  if  known. 

3.  Class  to  which  the  substance  belongs,  and  indication  of  degree  of 
activity  where  possible,  such  as  simple  substance,  weak  acid,  active  base,  salt, 
carbohydrate,  etc.      Terms  like  acid,  base,  and  salt  imply  certain  properties, 
which  need  not  be  given  in  detail. 

4.  Substances  with  which  combination  occurs,  such  as  the  metals  (ex- 
ceptions named),  the  non-metals  (exceptions  named),  water,  ammonia,  etc.     The 
class  to  which  the  resulting  compounds  belong  should  be  named. 

5.  Oxidizing  or  reducing  agent,  if  such,  with  illustrations  and  indication 
of  limitations. 

6.  Other  specific  chemical  reactions,  such  as  hydrolysis,  if  a  salt;  action 
of  chlorine,  if  a  hydrocarbon;  etc. 

Each  chemical  property  should  specify  directly  (or  by  implication,  e.g., 


490  INORGANIC  CHEMISTRY 

by  using  the  word  base)  some  definite  variety  of  chemical  change,  or  kind 
of  chemical  behavior,  and  should  name  the  materials  or  classes  of  materials 
involved. 

The  chemical  relations  are  not  properties  of  a  substance,  but  qualities  of  an 
element  in  combination,  such  as  atomic  weight,  valence,  non-metallic  or  metallic 
element  (oxide  acidic  or  basic,  halides  hydrolyzed  or  not).  They  have  been 
discussed  elsewhere  (pp.  226,  464). 

It  may  be  well  to  illustrate  inept  ways  of  giving  chemical  properties.  "Oxy- 
gen supports  combustion  of  a  candle."  The  beginner  does  not  know  that  a 
candle  is  composed  of  a  cotton  wick  surrounded  by  a  mixture  of  hydrocarbons 
and  fatty  acids,  so  no  chemical  reaction  can  be  extracted  by  him  from  this  state- 
ment. The  chemical  property  is  the  tendency  of  compounds  of  carbon  and 
hydrogen  to  interact  vigorously  with  oxygen,  giving  water  and  carbon  dioxide. 
The  burning  of  a  candle  for  light  is  an  application  of  this  property,  but  is  not  the 
property  itself.  Applications  should  be  given,  of  course,  but  they  should  follow 
the  property,  not  displace  it.  Similarly,  "bleaching  ability"  is  often  given  as  a 
chemical  property.  But  sodium  hyposulphite  and  other  substances  bleach  in- 
digo by  reducing  it  to  indigo  white,  and  hypochlorous  acid  and  ozone  bleach 
indigo  by  oxidizing  it  to  isatin,  sodium  hydroxide  bleaches  blue  prints  because  it 
is  an  active  base  and  interacts  with  the  ferrous  ferricyanide,  and  water  and  other 
solvents  sometimes  bleach  by  dissolving  the  dye.  Bleaching  may,  therefore,  be 
due  to  entirely  different  chemical  or  physical  properties  in  different  cases.  It  is 
an  application  of  some  specific  chemical  property,  and  should  be  given  as  an 
illustration  of  the  use  of  the  property.  Similarly,  "disinfecting  power"  in  the 
case  of  hypochlorous  acid  or  hydrogen  peroxide  is  due  to  oxidation  of  the  unstable 
substances  in  the  pathogenic  organisms,  but  sulphurous  acid  adds  itself  to  the 
aldehydes  which  their  protoplasm  contains,  alcohol  (used,  e.g.,  to  sterilize  the 
skin  before  vaccination)  takes  the  place  of  the  moisture  in  the  organisms,  killing 
them  by  physical  means,  and  formaldehyde  is  a  reducing  agent,  and  is  also  able " 
to  add  itself  to,  or  to  give  condensations  with,  many  organic  compounds.  Dis- 
infecting power  is  thus  an  application  of  one  of  the  foregoing,  or  of  some  other 
property,  which  should  first  be  specified.  When  a  substance  is  poisonous,  this 
should  be  emphasized,  but  not  given  as  a  chemical  property.  Even  when  the 
chemical  changes,  which  the  poisons  bring  about  in  the  body,  are  known,  they 
can  seldom  be  explained  in  inorganic  or  elementary  general  chemistry. 

Again  the  phrase  "supports  combustion"  is  of  uncertain  meaning.  In  ordi- 
nary language  it  refers  to  fuels  like  wood  and  coal.  To  say  that  "  chlorine  supports 
combustion,"  is  therefore,  not  specific  enough.  Coal  and  wood  will  not  burn  in 
it.  Iron,  copper,  and  antimony  will,  but  if  these  substances  are  not  named,  the 
statement  leaves  the  reader  to  supply  the  very  information  which  he  lacks,  and 
which  the  book  should  have  furnished.  Still  again,  the  statement  that  chlorine 
combines  with  copper  and  sodium  omits  to  make  it  clear  that  chlorine  combines 
easily  with  all  the  familiar  metals,  excepting  platinum  and  gold.  If  each  individ- 
ual element  with  which  chlorine  combines  is  to  constitute  a  separate  chemical 
property,  then,  with  only  two  elements  named,  the  list  of  this  class  of  properties 


OXIDATION  AND  REDUCTION  491 

is  very  incomplete,  since  it  includes  only  two  out  of  about  sixty-five.  Naming 
the  class  of  elements,  with  illustrations  and  exceptions,  covers  the  point,  and  puts 
less  strain  upon  the  memory. 

Negative  statements  convey  little  information.  "Nitrogen  will  not  support 
combustion."  This  means,  either  that  coal  and  tapers  will  not  burn  in  it,  in 
which  case  it  should  be  explained  that  carbon  and  hydrogen  do  not  readily  unite 
with  nitrogen,  or  else  that  nitrogen  is  not  oxygen!  Chemical  properties  should  be 
stated  positively,  not  reached  indirectly,  by  excluding  one  by  one  all  the  things 
the  substance  cannot  do.  Nitrogen  does  not  bleach,  does  not  burn,  and  can 
neither  climb  a  tree  nor  eat  grass  —  but  a  hundred  negative  items  will  still  leave 
us  ignorant  of  what  it  can  do.  When  two  elements  are  very  similar,  yet  one  has 
a  property  which  the  other  lacks,  it  is  instructive,  and  therefore  desirable,  to 
mention  the  absence  of  the  property,  but  such  cases  are  rare.  To  parade  as  a 
property  of  phosphorus  pentoxide  that  it  is  not  combustible,  is  like  saying  that  a 
bankrupt  cannot  pay  his  debts  —  it  is  redundant.  Even  more  absurd  is  the 
statement,  when  given  as  a  property,  that  chlorine  is  formed  by  oxidation  of 
hydrochloric  acid.  This  is  a  property  of  the  acid  and  oxidizing  agent.  Chlorine 
has  no  properties  as  a  substance  until  after  it  has  been  formed. 

OXIDATION  AND  REDUCTION 

Oxidation  by  Oxygen.  —  The  simplest  oxidations  are  the  cases 
where  a  metal  or  non-metal  unites  with  oxygen: 

2Cu  +  O2  ->  2CuO,        S  +  O2  ->  S02. 

Union  of  a  compound  with  additional  oxygen  is  oxidation  also. 
2SO2  +  O2  ->  2S03,        3KC1O  -» 2KC1  +  KC1O3. 

The  removal  of  hydrogen  from  hydrogen  chloride  (preparation  of 
chlorine,  p.  217),  is  also  denned  as  oxidation. 

O2  +  4HC1  ±=?  2H2O  +  201,. 
2KMn04  +  16HC1  -4  8H2O  +  2KC1  +  2MnCl2  +  501,. 

Every  oxidation  is  accompanied  by  reduction  of  the  oxidizing  agent. 
Thus,  in  the  second  last  equation,  the  free  oxygen  is  reduced  to 
water.  Again,  in  the  third  last  equation,  2KC1O  is  reduced  to 
2KC1,  while  1KC10  becomes  KC1O3  by  oxidation. 

In  the  laboratory,  we  frequently  discover  that  an  oxidation  has 
occurred  by  noticing  the  presence  of  a  product  of  reduction. 
Thus,  when  we  heat  carbon  with  sulphuric  acid :  2H2SO4  +  C  — > 
C02  +  2H20  +  2SO2,  we  do  not  notice  the  product  of  oxidation; 
CO2,  because  it  is  odorless  and  colorless,  but  we  perceive  at  once 
the  odor  of  the  sulphur  dioxide,  and  realize  that  the  sulphuric  acid 


492  INORGANIC  CHEMISTRY 

must  have  oxidized  some  substance,  for  this  gas  could  not  have 
been  formed  (at  the  temperature  employed)  except  by  reduction. 

Note  that  the  removal  of  the  elements  of  water  is  neither  oxida- 
tion nor  reduction,  for  equivalent  amounts  of  both  oxygen  and 
hydrogen  are  removed: 

2HC10  ->  H20  +  ClaO,        H2C03  ->  H2O  +  C02. 

In  all  the  cases  discussed  above,  oxidation  consists  in  adding  oxy- 
gen or  removing  hydrogen. 

Oxidation  by  Other  Negative  Elements.  —  Oxygen  is  only 
one  of  the  class  of  elements  called  non-metallic  or  negative  ele- 
ments, so  we  cannot  logically  restrict  the  term  "oxidation"  to 
actions  involving  oxygen.  Thus,  forming  a  chloride,  or  increasing 
the  proportion  of  chlorine  in  a  compound  is  oxidation : 

Cu  +  C12  ->  CuCl2,        2FeCl2  +  C12  '*•  2FeCl3. 

In  every  compound,  one  of  the  elements  is  relatively  positive  and 
the  other  relatively  negative.  Thus,  copper  is  positive  and 
chlorine  negative.  In  carbon  dioxide  C02,  carbon  is  (relatively) 
positive  and  oxygen  negative,  and  in  calcium  carbide  CaC2  cal- 
cium is  positive  and  carbon  (relatively)  negative. 

Thus,  oxidation  is  introducing,  or  increasing  the  proportion  of  the 
negative  element,  or  removing,  or  reducing  the  proportion  of  the 
positive  element.  Reduction  is  the  converse. 

Oxidation  and  Valence.  —  Combining  a  metal  with  oxygen 
or  sulphur  raises  the  active  valence  of  .the  metal  from  zero  to  some 
finite  value:  2Cu°  +  02°  -^  2CuIIO11.  Metallic  copper  has  no 
valence  in  use.  In  CuO  or  CuCl2  it  has  gained  the  valence  II. 
The  copper  has  been  oxidized.  Similarly,  changing  FeCl2  into 
Feds  increases  the  active  valence  of  the  iron  from  II  to  III  (oxida- 
tion). Conversely,  changing  2HC1  to  C12  decreases  the  active 
valence  of  chlorine  from  I  to  zero  (oxidation).  In  the  same  equa- 
tion (p.  218),  KMn  in  KMnO4  must  have  a  total  valence  of  VIII, 
but  in  the  products  KC1  +  MnCl2  the  total  valence  has  decreased 
to  III  (reduction). 

Again,  in  displacement,  e.g.,  Zn  -f  2HC1  — >  ZnCl2  +  H2,  the 
zinc  is  oxidized  because  the  active  valence  goes  from  zero  to  II, 
and  the  hydrogen  is  reduced. 

Hence,  oxidation  consists  in  increasing  the  active  valence  of  a 


OXIDATION  AND  REDUCTION  493 

positive  element  or  decreasing  that  of  a  negative  element.  Reduc- 
tion is  the  converse. 

This  way  of  stating  the  rule  makes  it  clear  why  removing  the 
elements  of  water  is  neither  oxidation  nor  reduction.  We  are  re- 
moving both  a  positive  and  a  negative  element,  and  are  removing 
them  in  equi-valent  amounts,  2H1  -J-  On. 

Oxidation   and  lonization.  —  If,  in  the  last  illustration,  we 

write  the  equation  ionically:  Zn  -I-  2H+  — >  Zn++  +  H2,  we  dis- 
cover that,  logically,  we  must  consider  the  change  from  metallic 
zinc  to  zinc-ion  to  be  in  itself  oxidation.  This  is  the  case  whether 
the  zinc-ion  later  combines  with  a  negative  ion  to  form  a  molecule 
or  not.  Mere  union  or  disunion  of  ;ons  is  neither  oxidation  nor 
reduction.  Conversely,  the  discharge  of  the  2H+  giving  H2  is  re- 
duction. 

Thus,  ionization  of  an  elementary  substance  to  form  a  positive 
ion  is  oxidation,  and  ionization  to  form  a  negative  ion  is  reduction, 
and  conversely. 

Oxidation  and  Electrons.  —  Increasing  the  valence  of  an 
atom  of  a  positive  element  (oxidation)  consists  in  removing  one  or 
more  electrons:  Na°  —  «  =  Na+  (p.  354).  Increasing  the  valence 
of  an  atom  of  a  negative  element  (reduction)  means  adding  one  or 
more  electrons :  Cl°  +  c  — >•  Cl~. 

Hence,  oxidation  is  removing  electrons  and  reduction  is  adding 
electrons. 

Making  Equations  for  Oxidations  and  Reductions.  —  The 

writing  of  equations  for  actions  involving  oxidation  and  reduction, 
where  there  are  more  than  two  substances  on  one  side  of  the  equation, 
is  difficult,  and  .a  system  or  plan  is  of  great  value.  The  plan  of 
partial  equations  (p.  270)  is  often  helpful.  There  are  three  other 
systems  which  are  in  use.  (1)  When  the  action  involves  oxygen 
acids  and  their  salts,  the  formulae  can  be  rewritten  so  as  to  show  the 
anhydride  (see  below).  (2)  The  second,  called  the  system  of  positive 
and  negative  valences,  is  more  generally  applicable  (next  section). 
(3)  The  third  describes  oxidation  in  terms  of  ions  and  positive  elec- 
trical charges  (p.  496). 

Making  Equations:    Using  Positive  and  Negative  Valences 

(p.  425).  —  1.   Each  compound  is  composed  of  elements  which  are, 


494  INORGANIC  CHEMISTRY 

relatively  to  one  another,  either  positive  or  negative.  Thus,  in  KMn04, 
K  and  Mn  are  positive  and  O  is  negative.  In  CS2,  C  is  (relatively) 
positive  and  S  negative.  We  say,  then,  that  C  has  a  positive 
valence  of  four  (+4)  and  S  has  a  negative  valence  of  two  (  —  2), 
just  as  it  has  in  H2S. 

2.  In  every  compound,  the  algebraic  sum  of  the  positive  and  negative 
valences  must  be  zero.     Thus,  in  CS2  the  sum  is  +4  —  2X2  =  0 
(C++S2=).     This  is  simply  the  rule  of  equi- valence  (p.  136),  with  the 
addition  of  the  idea  of  relative  positiveness  and  negativeness. 

This  enables  us  to  determine  the  valence  of  each  element  in  a 
compound  like  KMnO4.  K+  is  always  univalent  and  positive.  0=, 
in  inorganic  compounds,  is  always  bivalent  and  negative.  The 
valence  of  Mn  has  different  values:  MnMCl2,  Mn2mO3,  MnIVO2, 
Mn2VII07,  etc.  By  the  rule  (sum  of  valences  equals  zero)  we  can 
tell  the  valence  of  Mn  in  this  compound.  The  valence  of  O4  (40~) 
is  -8.  That  of  K  is  +1.  That  of  Mn  must  therefore  be  +7 
(KMn+VII04).*  Again,  in  HC103,  the  valence  of  03  is  -6,  that  of 
H  is  +1,  therefore  that  of  Cl  must  be  +5.  Still  again,  in  K2Cr207, 
the  valence  of  Oy  is  —14,  that  of  K2  is  +2,  that  of  Cr2  is  therefore 
+  12,  and  that  of  Cr  necessarily  +6  (K2Cr2+VIO7) . 

3.  Since  rule  2  applies  to  every  compound  used  or  produced  in 
a  chemical  change,  it  follows  that  when  in  a  reaction  the  valence  of 
an  element  changes  in  value,  that  of  one  or  more  of  the  other  elements 
must  also  Change,  so  as  to  maintain  the  equality  of  +  and  —  valences. 
Thus,  if  one  element  loses  in  valence,  to  the  extent  of  +6,  some 
other  element  (or  elements)  must  lose  —6,  or  gain  +6.     The  gain 
(or  loss)  of  one  element  must  cancel  the  gain  (or  loss)  of  some  other 
element. 

4.  The  valence  of  a  free  element,  that  is,  its  active  valence,  is 
zero.     A  free  element  is  also  neutral  —  neither  positive  nor  nega- 
tive —  because  it  is  not  combined  with  any  other  element. 

Illustration  of  rules  3  and  4.  Thus,  in  the  action  for  preparing 
chlorine  with  manganese  dioxide  (p.  219) : 

Mn02  +  4HC1  ->  MnCl2  +  2H20  +  C12, 

4H  (4H+)  has  the  valence  +4  on  both  sides.  On  the  left  side, 
4C1  (4C1~)  has  the  valence  —4:  on  the  right  2C1  has  the  valence 
—  2,  and  C12  has  the  valence  0.  So  far  as  chlorine  is  concerned, 

*  The  reader  should  write  this,  and  other  formulae  discussed  below,  so  as  to, 
show  the  valences  thus:  K+MnJJJ+Or  (cf.  p.  425), 


OXIDATION  AND  REDUCTION  495 

there  is  a  change  from  —4  to  —2,  or  a  difference  of  —2.  Again, 
on  the  right,  Mn  has  the  valence  +2,  while  on  the  left  side  it  has 
the  valence  +4,  a  difference  of  +2.  The  two  differences,  —2 
and  +2,  cancel  one  another.  Stated  otherwise,  manganese  lost 
+2  and  chlorine  lost  —2,  so  that  the  other  +  and  —  valences  still 
in  use  remained  equal  in  number,  and  equi-valence  was  preserved. 

Balancing  an  Equation.  Suppose  we  wish  to  balance  the  equa- 
tion for  the  decomposition  of  chloric  acid  HC1O3.  We  ascertain, 
in  the  laboratory,  that  the  products  are  perchloric  acid  HC1O4, 
chlorine  dioxide  C102,  and  water. 

Skeleton:  HC1O3  ->  HC1O4  +  C102  +  H2O. 

Since  H+  and  0~  do  not  change  in  valence,  only  Cl  has  been  affected. 
On  the  left  side,  the  valences  are  03  =  —6,  H  =  +1,  Cl  there- 
fore =  +5.*  On  the  right  side,  in  HC1O4,  the  total  valence  of 
oxygen  is  —8  and  of  hydrogen  +1.  That  of  Cl  is  therefore  +7. 
In  C1O2,  the  valence  of  O2  is  —4,  and  that  of  Cl  therefore  +4.  Thus, 
Cl  changes,  from  +5,  partly  to  +7  and  partly  to  +4.  To  achieve 
this,  arithmetically,  we  require  3C1  on  the  left  (=  3  X  +5  =  +15), 
giving  Cl  =  +7  and  2C1  =  2  X  +4  =  +8,  or  a  total  of  +15  on 
the  right.  Thus,  we  require  3HC103: 

Balanced:  3HC103  =  HC104  +  2C102  +  H20. 

Balancing  Another  Equation.  In  the  reaction  for  preparing 
chlorine  (p.  218),  the  skeleton  is: 

Skeleton:    KMnO4  +  HC1  ^4  H20  +  KC1  +  MnCl2  +  Cl. 

Here,  in  KMn04,  the  valence  of  Mn  is  +7.  In  MnCl2  it  is  +2, 
a  loss  of  +5.  The  chlorine  also  changes  its  valence  from  —1  to  0, 
a  loss  of  —1.  Evidently,  so  that  the  changes  may  cancel  out,  for 
every  Mn  losing  +5,  5C1  must  lose  5  X  —  1  and  be  liberated: 

Incomplete:    KMnO4  +  HC1  -»  H2O  +  KC1  +  MnCl2  +  5C1. 

Since  there  is  now,  altogether,  8C1  on  the  right,  8HC1  will  be  re- 
quired on  the  left.  The  8H  will'  give  4H2O : 

Balanced:        KMnO4  +  8HC1  -»  4H2O  +  KC1  +  MnCl2  +  5C1. 
Molecular:  2KMn04  +  16HC1  ->  8H2O  +  2&C1  +  2MnCl2  +  5C12. 

For  another  method  of  balancing  this  equation,  see  p.  496. 
*  Write  these  (and  other  formula)  thus:  H+C4^+O3=  (cf.  p.  425). 


496  INORGANIC  CHEMISTRY 

Making  Equations,  by  Using  the  Anhydrides^  —  To  balance 
the  equation  for  the  decomposition  of  chloric  acid,  we  first  write 
the  skeleton  equation: 

Skeleton:  HC103  -»  HC104  +  C102  +  H2O. 

Then  we  divide  the  acids  into  water  and  the  anhydrides  (p.  484). 

Analyzed:        H20,C12O5  ->  H20,C12O7  +  C102  +  H20. 

We  now  perceive  that,  disregarding  the  water,  some  C12O5  must 
lose  oxygen  to  give  2C1O2  +  O,  and  that  some  C12O5  must  gain  2O, 
becoming  C12O7.  To  furnish  the  20,  clearly  2C12O5  is  required, 
giving  4C1O2  +  2O,  and  a  third  C1205  gains  this  20.  Thus,  alto- 
gether 3C12O5  will  be  required: 

Balanced:       3H2O,C12O5  ->  H2O,C1207  +  4C102  +  2H20 
or  6HC103  -*  2HC104  +  4C1O2  +  2H20. 

This  equation  is  then  divided  by  two  throughout. 

Making  Equations  by  Oxidation  of  Ions,  Using  Positive 
Electrical  Charges.  —  All  oxidation  reactions  involving  ionogens 
can  be  written  in  terms  of  ions.  Thus,  the  oxidation  of  hydro- 
chloric acid  by  potassium  permanganate  can  be  so  written.  The 
potassium-ion  clearly  is  not  affected,  and  may  be  omitted.  The 
ions  concerned  are: 

Mn04~  +  H+  +  Cr  -»  H20  +  Mn++  +  Cl°. 

Cl°  with  no  charge  stands  for  free  chlorine.  Now  we  can  divide 
the  action  into  (1)  the  behavior  of  the  oxidizing  agent,  which  is 
general,  and  will  be  used  wherever  the  same  oxidizing  agent  is  used; 
(2)  the  fate  of  the  substance  being  oxidized,  which  again  is  general, 
because  other  oxidizing  agents  will  change  it  in  the  same  way. 

Mn04-  +  8H+  -»  4H20  +  Mn++  +  50 .  (1) 

In  words,  each  permanganate  ion,  with  a  free  acid  present  (oxi- 
dizing mixture),  will  give  water,  manganous-ion,  and  a  balance  of 
five  unit  positive  charges. 

50  +5Cr-»5Cl°.  (2) 

100  +  50*=  -*  5H2O  +  5O2°.  (21) 

100  +  5S03=  +  5H2O  ->  5SO4=  +  10H+.  (211) 

These  three  equations  represent  the  oxidation  of  (2)  hydrochloric 
acid,  or  (21)  hydrogen  peroxide,  giving  free  oxygen,  or  (2n)  sul- 


OXIDATION  AND  REDUCTION  497 

phurous  acid,  with  water  furnishing  the  oxygen,  and  leaving  the 
solution  strongly  acid  (=  5H2SO4).     Note  that  the  sums  of  the  + 
and  —  charges  on  opposite  sides  of  each  equation  are  equal. 
To  obtain  the  final  ionic  equation,  add  (1)  and  (2) : 

Mn04~  +  8H+  -» 4H20  +  Mn++  +  58 .  (1) 

50  +5Cr->5Cl°. (2) 

MnOr  +  8H+  +  5C1-  -»  4H2O  +  Mn++  +  501°. 

Before  adding  (1)  and  (21)  and  (1)  and  (2XI),  the  first  equation  (1) 
must  be  doubled  throughout,  so  that  the  10©  may  cancel  out. 

Exercises.  —  1.  Assign  to  its  proper  class  (pp.  228,  402)  each  of 
the  actions  mentioned  in  this  chapter. 

2.  Knowing  that  sodium  acid  tartrate  NaHC4H4O6  is  insoluble 
in  50  per  cent  alcohol,  how  should  you  make  chloric  acid  (p.  482)  ? 

3.  Make  the  equation  for  the  interaction  of  chlorine  with  calcium 
hydroxide  in  hot  water  (p.  480).     How  should  you  make  zinc  chlorate 
from  zinc  hydroxide  Zn(OH)2? 

4.  How  should  you  make  pure  potassium  hypochlorite  from 
hypochlorous  acid  (p.  474)? 

5.  Explain,   in  terms  of  ionic  equilibrium,   why  dilute  hypo- 
chlorous  acid  can  be  obtained  by  adding  one-half  of  an  equivalent 
of  an  active  acid  (p.  475)  to  bleaching  powder,  and  distilling  the 
mixture. 

6.  On   what   circumstances   would   the   possibility   of   making 
barium  chlorate  by  action  of  chlorine  on  barium  hydroxide  depend 
(p.  481)?     Could  pure  barium  chlorate  be  obtained  easily  by  this 
means  (see  Table  of  Solubilities)? 

7.  Make  the  equations  for:    (a)  the  preparation  of  potassium 
bromate;     (6)   pure  aqueous  bromic  acid;     (c)   the  interaction  of 
iodine  with  aqueous  potassium  hydroxide  in  the  cold,  and  (d)  when 
heated. 

8.  Make  the  equations  for  the  interactions  of  chlorine  dioxide 
with  water,  and  with  aqueous  potassium  hydroxide. 

9.  Find  the  formulae  of  the  anhydrides  of  the  following  acids: 

HP03,  H2Se04,  H3As03,  H3As04,  H6S06. 

10.  Find  the  formulae  of  the  anhydrides  of  the  acids  from  the 
following  formulae  of  salts: 

a,  NaJHPO*,  NaH2P03, 


498  INORGANIC  CHEMISTRY 

11.  Classify  the  following  changes  as  oxidations  or  reductions, 
(a)  H2Cr207-»H2Cr04  +  Cr03;  (6)  HMnO4-*Mn02;  (c)  P->I"; 
(d)  2H202  ->  2H2O  +  02. 

12.  Using  positive  and  negative  valences,   determine  whether 
each  of  the  following  formulae  is  correct  or  incorrect:    Ca(Mn04)2, 
A1(C104)3,  Na*HIOB. 

13.  Apply  each  of  the  three  methods  (pp.  493,  496)  of  writing 
equations  to  the  four  following  reactions:     (a)   chlorine-water  on 
bromine;     (6)    chlorine-water   on    hydrogen   sulphide,    giving   free 
sulphur;    (c)  potassium  permanganate  and  free  acid  on  hydrogen 
sulphide,  giving  free  sulphur;    (d)  potassium  dichromate  and  free 
acid  (p.  418)  on  hydrogen  sulphide,  giving  the  chromic  salt  of  the 
acid  (Crm)  and  free  sulphur. 

14.  Explain  why  chlorine  is  much  less  soluble  in  a  solution  of 
sodium  chloride  than  in  water. 

15.  Why  does  light  liberate  oxygen  from  chlorine-water,  but  not 
from  chlorine  hydrate? 

16.  Why  does  a  given  weight  of  chlorine  in  the  form  of  hypo- 
chlorous  acid  have  twice  as  great  a  bleaching  (or  oxidizing)  capacity 
as  has  the  same  weight  of  chlorine  in  chlorine-water? 

17.  What  would  be  the  result  of  leading  chlorine  monoxide  gas 
into  concentrated  hydrochloric  acid? 

18.  When  bromic  acid  acts  upon  iodine  (p.  486),  why  is  HBr, 
and  not  Br2,  formed? 


CHAPTER  XXIV 
THE  ATMOSPHERE.     THE  HELIUM  FAMILY 

WE  have  seen  that,  to  counterbalance  the  pressure  of  the  air  over 
a  certain  area,  a  column  of  mercury  of  the  same  diameter  averaging 
760  mm.  in  height  is  required.  Let  the  section  of  the  column  be 
1  sq.  cm.  Then  the  pressure  of  a  column  of  air  1  sq.  cm.  in  section, 
and  extending  so  far  from  the  earth  as  any  downward  tendency  of  the 
air  exists,  is  equal  to  the  weight  of  a  column  of  mercury  containing 
76  c.c.  of  the  metal.  The  weight  of  1  c.c.  of  mercury  being  13.6  g., 
this  volume  of  the  metal  weighs  1033.6  g.  This  number  represents 
therefore  the  pressure  which  is  exerted  by  the  air  upon  each  square 
centimeter  of  the  earth's  surface.  In  ordinary  units  of  measure,  this 
is  nearly  fifteen  pounds  to  the  square  inch. 

A  more  vivid  appreciation  of  the  reality  of  this  pressure  may  be  obtained  by 
noticing  one  of  its  effects.  By  boiling  a  small  quantity  of  water  in  a  tin  can 
furnished  with  a  narrow  opening,  we  remove  the  whole  of  the  air  from  its  interior, 
displacing  it  by  steam.  While  the  boiling  is  in  progress,  we  suddenly  close  the 
opening  with  a  tightly  fitting  cork  and  remove  the  burner.  While  the  steam  was 
still  issuing  from  the  opening,  its  pressure  was  practically  that  of  the  atmosphere, 
and  the  can  was  subject  to  the  same  pressure  inside  and  out.  With  the  removal 
of  the  flame,  however,  the  steam  condenses,  and  the  pressure  on  the  interior  is 
reduced  to  a  minute  fraction  of  its  original  value,  while  the  pressure  on  the  ex- 
terior is  still  the  same  (1  atmosphere).  Under  this  pressure  a  vessel  of  ordinary 
tin-plate  completely  collapses  [Lect.  exp.]. 

Components  of  the  Atmosphere.  —  There  are  three  classes 
of  components  in  the  air.  Those  of  the  first  class,  oxygen,  nitrogen, 
and  the  inert  gases  of  the  helium  family,  are  present  in  almost  con- 
stant proportions.  Those  of  the  second  class  are  very  variable  in 
quantity,  and  include  carbon  dioxide,  water  vapor,  and  dust.  Those 
of  the  third  class,  such  as  the  sulphur  dioxide  in  city  air,  are  accidental. 

Components  which  are  Constant  in  Amount.  —  The  deter- 
mination of  the  oxygen  by  burning  phosphorus  in  air,  and  measuring 
the  residual  gas,  is  not  capable  of  application  in  an  exact  manner.  A 

499 


500 


INORGANIC  CHEMISTRY 


FIG.  115. 


stick  of  phosphorus,  enclosed  in  wire  gauze  (Fig.  115),  removes  the 
oxygen  rather  slowly  [Lect.  exp.].  It  is  better  to  use  a  large  amount 
of  phosphorus  in  the  form  of  thin  wire.  In  this  way  a  great  surface  is 
obtained,  and  the  absorption  of  oxygen  from  a  sample  of  air  may  be 
carried  out  in  a  few  seconds.  This  method  gives  fairly 
accurate  results,  since  there  is  no  time  for  any  appreciable 
change  in  the  temperature  or  pressure  of  the  atmosphere 
during  the  experiment.  Several  oxygen  acids  of  phos- 
phorus are  formed.  The  passage  of  purified  air  over 
heated  copper  has  also  been  used  for  the  same  purpose. 
When  this  method  is  employed,  the  volume  of  the  nitro- 
gen and  argon  which  survives  the  action  of  the  copper  is 
measured,  while  the  increase  in  weight  of  the  copper, 
through  formation  of  cupric  oxide,  gives  the  weight  of 
the  oxygen  which  was  originally  mixed  with  it. 

Still  another  method,  which  is  in  constant  employ- 
ment in  the  analysis  of  mixtures  of  gases,  may  also  be 
applied  to  the  air.  It  consists  in  bringing  a  measured 
volume  of  air  in  contact  with  an  alkaline  solution  of 
potassium  pyrogallate,  which  quickly  absorbs  the  oxygen, 
and  noting  the  decrease  in  volume. 

In  the  air  taken  from  mines,  from  mountain  tops,  from  the  surface 
of  the  sea,  and  from  inland  regions,  the  proportion  of  oxygen  to  the 
residual  gas  is  found  to  be  fairly  constant,  although  easily  perceptible 
differences  are  noted.  The  percentage  of  oxygen  in  dried  air  ranges 
between  20.26  and  21.00,  the  latter  being  the  proportion  in  normal 
air. 

When  the  residual  gas  is  led  slowly  through  a  heated  tube  con- 
taining magnesium,  the  nitrogen  unites  with  the  metal  to  form  the 
solid  nitride  Mg3N2,  and  only  about  10  c.c.  out  of  every  liter  remains 
uncombined.  This  residuum  is  argon,  mixed  with  0.15  per  cent  of 
its  volume  of  other  gases  belonging  to  the  helium  family  (see  below). 
Exact  measurement  by  volume  gives  78.06  per  cent  of  nitrogen  and 
0.94  per  cent  of  argon  in  dried  air. 

Gaseous  Components  which  are  Variable  in  Amount.  — 

Pure  country  air  contains  about  3  parts  in  10,000  of  carbon  dioxide. 
In  city  air  there  are  from  6  to  7  parts  in  the  same  volume,  while  in  the 
air  of  audience-rooms,  where  the  ventilation  is  defective,  the  propor- 
tion may  rise  as  high  as  50  parts. 

The  simplest  way  of  showing  the  presence  of  carbon  dioxide  in  the 


THE  ATMOSPHERE.      THE  HELIUM   FAMILY  501 

air  is  by  exposing  a  solution  of  barium  hydroxide,  an  active  base,  in  a 
shallow  vessel.  After  a  short  time  a  layer  of  barium  carbonate  forms 
upon  the  surface:  Ba(OH)2  +  CO2  -*  BaCO3  j  +  H20.  The  same 
action  may  be  utilized  for  the  purpose  of  quantitative  analysis.  A 
measured  volume  of  air  is  bubbled  slowly  through  a  measured  volume 
of  a  solution  of  barium  hydroxide  of  known  concentration,  and  the 
quantity  of  barium  hydroxide  remaining  is  determined  by  titration 
(p.  390). 

The  sources  of  the  carbon  dioxide  in  the  air  are  numerous.  It 
comes  from  the  decay  of  vegetable  and  animal  matter,  in  which, 
chiefly  through  the  influence  of  minute  vegetable  organisms,  the 
carbon  is  oxidized  to  carbon  dioxide.  It  is  formed  also  by  the  com- 
bustion of  coal  and  wood,  but  the  thirteen  hundred  million  tons 
of  coal  burned  annually,  giving  three  times  that  weight  of  carbon 
dioxide,  would  add  only  one-six  hundredth  to  the  total  present  in 
the  air.  It  is  exhaled  by  animals,  being  produced  in  the  body  by 
oxidation  of  the  carbon  in  the  food  which  they  eat.  It  also  issues 
from  the  earth,  in  volcanic  as  well  as  in  other  neighborhoods.  The 
proportion  of  this  gas  in  the  air  would  naturally  increase  continuously, 
though  slowly,  as  the  result  of  these  processes,  were  it  not  that  it  is 
removed  just  as  continuously  by  the  action  of  growing  plants  (see 
p.  579),  which  use  it  as  food.  It  may  be  added,  also,  that  carbon 
dioxide,  being  a  soluble  gas,  is  contained  in  sea  water,  dissolved  and 
as  Ca(HCO3)2,  and  the  total  amount  in  the  ocean  is  much  greater  than 
that  in  the  air.  The  removal  by  plants  and  by  solution  in  sea  water 
thus  keeps  the  proportion  in  the  air  fairly  constant. 

The  presence  of  carbon  dioxide  in  the  breath  may  be  shown 
very  quickly  by  blowing  through  a  tube  into  calcium  hydroxide 
solution  (limewater).  Calcium  carbonate  CaCOs  is  precipitated. 
We  draw  about  500  c.c.  of  air  into  our  lungs  at  each  breath,  or  half 
a  cubic  meter  per  hour.  In  the  lungs,  some  oxygen  is  removed,  the 
percentage  by  volume  falling  from  21  to  16,  and  we  add  some  carbon 
dioxide,  the  proportion  increasing  from  0.03  in  country  air  to  about 
4  per  cent.  A  candle  flame  is  extinguished  in  exhaled  air,  because 
the  maintenance  of  such  a  flame  requires  at  least  18.5  per  cent  of 
oxygen.  But  air  will  sustain  life  until  the  proportion  has  fallen  to 
about  10  per  cent. 

The  proportion  of  water  vapor  is  constantly  changing.  When 
the  air  becomes  cool,  as  it  does  most  often  in  the  upper  layers,  the 
vapor  condenses  to  droplets,  forming  fogs  and  clouds.  When  the 
condensation  continues,  the  drops  become  larger  and  fall  as  rain.  On 


502  INORGANIC  CHEMISTRY 

the  other  hand,  when  the  weather  is  warm,  water  from  the  soil,  and 
from  rivers,  lakes,  and  oceans,  passes  into  vapor  and  the  amount  in 
the  air  increases. 

The  ammonium  nitrate  rises  from  the  interaction  of  nitric  acid 
and  ammonia.  The  latter  is  formed  by  the  decay  of  animal  matter 
(p.  517);  the  former  by  the  union  of  nitrogen  and  oxygen  during 
thunder-storms.  The  electrical  discharges  produce  nitrogen  tetroxide 
(see  p.  533),  which  with  water  gives  nitric  acid  (q.v.). 

Humidity.  —  The  moisture  in  the  air  is  usually  defined  in  terms 
of  the  relative  humidity,  the  standard  being  the  quantity  required 
to  saturate  the  air.  The  open  air  is  seldom  actually  saturated,  but, 
when  a  portion  is  confined  in  a  vessel  over  water,  it  soon  becomes 
so.  The  humidity  is  then  100  per  cent.  If  the  partial  pressure  of 
water  vapor  present  is  only  half  as  great  as  the  vapor  pressure  of 
water  at  the  same  temperature,  the  humidity  is  50  per  cent  The 
average  humidity  of  the  atmosphere  is  roughly  about  66  per  cent. 

At  18°  (64.4°  F.),  the  vapor  pressure  of  water  is  15.4  mm.  Thus 
air  saturated  with  moisture  at  18°  (100  per  cent  humidity)  would 
contain  15.4/760,  or  about  2  per  cent  by  volume  of  water  vapor. 
If  this  air  were  cooled  to  0°  (32°  F.),  a  temperature  at  which  the 
vapor  pressure  of  water  is  only  4.6  mm.,  the  air  could  retain  only 
4.6/760,  or  0.6  per  cent,  of  moisture.  The  difference,  amounting 
to  10.4  g.  (10.4  c.c.)  of  water  per  cubic  meter,  would  condense  as 
fog  or  rain. 

The  proportion  of  water  in  a  given  volume  of  air  may  be  meas- 
ured most  accurately  by  permitting  the  air  to  stream  slowly  through 
tubes  filled  with  calcium  chloride  or  phosphoric  anhydride.  The 
increase  in  weight  of  the  charged  tubes  represents  the  quantity 
of  moisture  abstracted  from  the  sample.  It  may  also  be  ascer- 
tained by  noting  the  temperature  to  which  air  has  to  be  cooled 
before  it  becomes  saturated  and  deposits  dew  (dew-point).  For 
example,  if  air  at  18°  has  to  be  cooled  to  11°  before  it  deposits  dew, 
it  contains  water  vapor  at  a  pressure  of  9.8. mm.  (Appendix  IV). 
If  saturated  at  18°,  it  would  have  contained  water  vapor  at  a  partial 
pressure  of  15.4  mm.  The  relative  humidity  was,  therefore,  9.8/15.4, 
or  63.6  per  cent. 

Ventilation.  —  On  a  moist  day,  we  speak  of  the  atmosphere 
as  " heavy"  or  "oppressive."  The  barometer,  however,  is  lower 
on  such  days,  and  the  pressure  below  the  average.  Moist  air  must 


THE  ATMOSPHERE.      THE  HELIUM  FAMILY  503 

be  lighter  than  dry  air,  because  in  moist  air  molecules  of  relative 
weight  18  (H2O)  have  been  substituted  for  an  equal  number  of  mole- 
cules of  oxygen  and  nitrogen  with  the  relative  weights  32  and  28. 
The  discomfort  is  due  to  a  different  cause. 

The  oxidation  of  digested  food  carried  by  the  blood  is  accom- 
panied by  liberation  of  heat,  yet  our  bodies  must  remain  at  98.6°  F. 
(37°  C.).  A  rise  of  a  few  tenths  of  a  degree  produces  discomfort. 
A  little  of  the  heat  is  lost  by  radiation  from  the  surface  of  the  body, 
but  the  real  adjustment  is  secured  by  evaporation  of  water  through 
the  skin.  The  vaporization  of  1  g.  of  water  (at  100°)  removes  heat 
amounting  to  540  calories  (603  cal.  at  37°  C.).  Evaporation  of  a 
single  ounce  (28 J  g.)  of  water  will  therefore  lower  the  temperature  of 
96.5  kilograms  (168  Ibs.)  of  water  (or  flesh,  which  is  largely  water)  by 
two-tenths  of  a  degree  C.  (nearly  0.4°  F.). 

The  " oppressive"  feeling,  then,  is  due  to  the  fact  that  the  air 
is  too  nearly  saturated,  evaporation  is  being  hindered  (p.  147),  and 
heat  is  accumulating.  Hence,  the  relative  humidity  is  the  measure 
of  the  goodness  or  badness  of  the  air  of  a  room. 

In  winter,  cold  and  therefore  relatively  dry  air  is  brought  into 
the  house  and  heated.  This  makes  the  relative  humidity  very  low, 
evaporation  proceeds  too  fast,  and  discomfort  follows.  In  summer, 
however,  the  outside  air  is  often  already  nearly  saturated  at  the 
temperature  of  the  room.  Unless  there  is  a  rapid  change  of  air  by 
ventilation,  the  moisture  from  the  bodies  of  those  in  the  room  in- 
creases the  humidity,  and  discomfort  arises  from  a  cause  opposite  to 
the  one  which  produced  it  in  winter. 

It  should  be  noted,  also,  that  even  though  the  air  is  in  constant 
motion,  the  layer  of  air  next  our  skin  (even  the  exposed  parts)  is 
hindered  from  moving  by  friction.  There  is  a  stationary  layer 
close  to  the  surface,  which  quickly  reaches  the  temperature  of  the 
body  and  becomes  saturated  at  that  temperature.  The  water 
molecules  can  leave  this  layer,  and  make  room  for  others,  only  by 
diffusion,  which  is  a  deliberate  rather  than  a  speedy  process.  Now, 
an  electric  fan,  although  it  brings  no  fresh,  dryer  air  into  the  room, 
nevertheless  stirs  the  air  and  blows  away  the  moist,  saturated  layer 
next  the  skin.  It,  at  least,  makes  this  layer  much  thinner,  and 
reduces  greatly  the  distance  the  water  molecules  have  to  go  by  mere 
diffusion.* 

*  The  same  conception  applies  to  dissolving  a  salt.  A  stationary  layer  of 
saturated  solution  is  formed  on  the  surface,  and  the  molecules  of  the  salt  can 
escape,  and  make  room  for  more,  only  by  diffusion.  In  liquids,  this  is  a  very 


504  INORGANIC  CHEMISTRY 

Formerly,  the  accumulation  of  carbon  dioxide  from  the  breath 
was  blamed  for  the  unhealthiness  of  unventilated  rooms.  The 
proportion  found  in  such  rooms,  however,  is  almost  never  sufficient 
to  do  any  harm.  Then,  it  was  imagined  that  traces  of  highly  poison- 
ous compounds  were  exhaled  by  the  body.  No  one,  however,  has 
yet  been  able  to  prove  that  such  poisons  exist. 

The  aims  of  ventilation  are,  therefore,  to  supply  fresh  outside 
air,  to  keep  it  in  motion,  and  to  maintain  a  humiolity  chat  is  neither 
too  low  nor  too  high. 

Dust  in  the  Air.  —  A  beam  of  sunlight,  crossing  a  dark  room, 
can  be  seen  by  the  light  reflected  from  the  particles  of  dust  in  the 
air.  The  dust  varies  both  in  kind  and  quantity  according  to  the 
locality.  It  is  found  to  be  partly  inorganic,  and  to  consist  of  clay, 
limestone,  and  soot  from  ill-burned  fuel.  In  factories  the  dust  may 
consist  of  minute  particles  of  glass,  steel,  cement,  or  other  substances. 
The  organic  dust  may  be  divided  into  two  kinds.  The  part  which  is 
dead  includes  coal  dust,  refuse  from  the  streets,  minute  shreds  of 
cotton,  linen,  hay,  etc.  The  living  dust  consists  of  pollen  grains, 
spores  of  fungi  and  molds,  bacteria,  and  similar  microscopic  organisms. 
The  presence  of  such  germs  in  the  air  is  shown  by  the  fact  that,  when 
nutritive  liquids  have  been  exposed  to  the  air,  even  for  a  few  minutes, 
putrefaction  very  soon  sets  in.  Some  germs  also  produce  disease 
when  they  land  on  a  place  where  the  skin  has  been  damaged  by  a  cut 
or  burn,  or  on  an  incision  made  in  the  course  of  an  operation.  After 
infection,  antiseptic  treatment,  e.g.,  with  hydrogen  peroxide,  destroys 
the  organisms.  But  protection  in  advance,  e.g.,  with  petrolatum  (q. 
v.),  until  a  new  skin  has  formed,  is  better. 

It  is  worth  noting  that  natural  soil  contains  about  100,000  micro- 
organisms per  c.c.,  good,  unfiltered  river  water  from  6000  to  20,000 
per  c.c.,  and  pure  air  only  4  or  5  per  liter. 

Flasks  can  be  filled  with  dustless  air  through  the  displacement  of 
that  which  they  contain  by  air  drawn  through  a  wide  tube  packed 
with  12-15  inches  of  cotton.  It  has  been  shown  by  Aitken  that  air 
filtered  in  this  way  behaves  differently  from  ordinary  air  in  respect  to 
the  manner  in  which  its  moisture  condenses. 

If  a  sample  of  moist  air  is  cooled  until  it  contains  more  water  vapor 
than  it  could  take  up  at  the  existing  temperature,  the  excess  of  mois- 

slow  process.  By  shaking  the  solid  and  liquid,  however,  the  stationary  layer 
is  partly  washed  away.  It  is  made  thinner,  so  that  the  distance  the  molecules 
have  to  travel  by  diffusion  is  greatly  reduced,  and  the  whole  operation  is  hastened. 


THE  ATMOSPHERE.      THE  HELIUM   FAMILY 


505 


ture  is  deposited.  This  deposition  usually  takes  place  by  the  forma- 
tion of  a  multitude  of  little  particles  of  liquid  water,  which  together 
make  up  a  fog.  Now  dustless  air  lacks  this  property  entirely.  When 
saturated  with  water  and  then  cooled,  it  does  not  give  any  trace  of  fog. 
The  excess  of  moisture  is  gradually  deposited  upon  the  walls  of  the 
vessel  and  upon  any  material  objects  which  it  contains,  but  of  fog 
there  is  no  trace  visible.  It  seems  that  the  particles  of  dust  are  re- 
quired as  nuclei  round  which  the  water  may  gather.  In  the  absence 
of  dust,  and  therefore  of  proper  nuclei,  the  moisture  is  not  precipi- 
tated in  the  usual  way.  Thus  fogs  and  rain  would  be  impossible  but 
for  the  presence  of  dust  in  all  ordinary  air.  In  the  absence  of  dust, 
the  cooling  would  produce  supersatura- 
tion,  which  would  be  slowly  relieved  by 
condensation  on  the  surfaces  of  houses, 
plants,  animals,  and  land.  Thus,  in  a 
dustless  atmosphere  an  awning  or  um- 
brella would  afford  no  shelter. 

The  formation  of  fog  in  ordinary 
air,  and  its  absence  in  filtered  air,  is 
easily  shown  in  a  darkened  room  (Fig. 
116).  The  flask  contains  some  water  to 
saturate  the  air.  When  suction  is  ap- 
plied, by  the  mouth,  to  the  tube  S,  the 
saturated  air  in  the  flask  expands  and  is  cooled.*  With  ordinary 
air,  a  fog,  brilliantly  illuminated  by  the  beam  of  light,  is  instantly 
oroduced.  Filtered  air  (dustless)  gives  no  fog.  On  the  other  hand, 
a  whiff  of  smoke  from  smoldering  paper,  when  admitted  to  the 
flask,  causes  a  fog  (after  cooling)  of  extraordinary  denseness  [Lect. 
exp.]. 

By  diluting  air  with  dustless  air,  generating  fog  in  the  mixture, 
and,  with  the  help  of  a  microscope,  counting  the  globules  when  they 
settle,  an  estimate  of  the  numb'er  of  particles  of  dust  in  air  may  be 
made.  It  is  found  that  rain  removes  a  large  proportion  of  them, 
while  respiration  and  combustion  greatly  increase  their  number. 
The  prevalence  of  fogs  in  cities  is  thus  accounted  for.  The  following 
are  the  numbers  of  dust  particles  in  1  c.c.  of  air: 

Outside,  raining 32,000     A  room,  near  the  ceiling    .     5,420,000 

Outside,  fair       130,000     Air  above  Bunsen  flame     .  30,000,000 

A  room        1,860,000 

*  Compression  with  a  bicycle  pump  heats  air,  and  expansion  cools  it. 


FIG.  116. 


506  INORGANIC  CHEMISTRY 

Air  a  Mixture.  —  Since  the  main  components  of  air  were  not 
definitely  identified  until  the  end  of  the  eighteenth  century,  we  can 
understand  why  the  substance  was  for  long  considered  to  be  an  ele- 
ment. The  experiments  which  we  have  described,  in  which  the 
oxygen  was  removed  from  the  air  and  the  nitrogen  remained,  do  not 
prove  that  the  original  constituents  were  present  simply  in  mechanical 
mixture.  They  might  have  been  combined,  and  the  combustion  of 
phosphorus,  for  example,  might  have  represented  the  removal  of 
oxygen  from  combination  with  nitrogen  and  its  appropriation  by  the 
phosphorus.  It  may  be  well,  therefore,  to  point  out  some  reasons 
which  lead  us  to  regard  the  air  as  a  mixture: 

1.  When  two  substances  enter  into  chemical  combination,  the  new 
body  invariably  has  different  physical  properties  from  either  of  the 
original  ones.     Each  of  the  substances  in  air,  however,  has  precisely 
the  same  properties  which  it  exhibits  when  free,  separate,  and  pure. 
This  is  characteristic  of  a  mixture.     Thus,  there  is  no  simple  relation 
between  the  refractive  power  for  light  which  a  compound  possesses  and 
the  refractive  powers  of  its  constituents.     In  the  case  of  air,  however, 
the  refractive  power  is  exactly  that  which  we  should  calculate  from 
the  refractive  powers  of  the  constituents,  taking  into  account  the 
proportions  of  them  which  air  contains. 

Again,  the  nitrogen  and  oxygen  dissolve  independently  in  water 
in  proportion  to  their  solubilities  and  partial  pressures  (p.  189).  If 
the  air  were  a  compound,  it  would  dissolve  as  a  whole,  and  the  rela- 
tive proportions  of  the  components  would  not  be  changed  by  the 
process.  Also,  the  density  of  air  is  precisely  that  which  we  find  by 
calculation  from  the  known  proportions  and  several  densities  of  the 
components.  Likewise,  when  liquefied  air  is  allowed  to  evaporate  in 
a  suitable  apparatus,  the  nitrogen,  being  more  volatile,  can  be  sepa- 
rated from  the  oxygen.  When  the  oxygen,  in  turn,  is  allowed  to 
evaporate,  the  carbon  dioxide  and  water  remain  as  solids,  frozen  at 
this  low  temperature.  No  compound  of  nitrogen  and  oxygen  is  found. 

2.  The  proportion  by  volume  in  which  the  gases  are  found  in  the 
air  is  not  so  simple  as  the  proportions  which  we  observe  in  cases  of 
chemical  combination.     The  proportion  is  close  to  4:1,  but  not 
exactly  4:1.     Besides,  as  we  have  seen,  the  proportion  is  not  per- 
fectly constant. 

3.  The  composition  of  air  varies,  while  the  composition  of  definite 
chemical  substances  is  always  the  same.     The  proportions  by  weight 
also  in  which  the  components  are  contained  in  air  are  not  integral 
multiples  of  the  atomic  weights  (see  Exercise  2). 


THE  ATMOSPHERE.      THE  HELIUM   FAMILY  507 

Composition  of  Air.  —  Air,  when  freed  from  carbon  dioxide 
and  water,  contains  by  volume  78.06  per  cent  of  nitrogen,  21.00 
per  cent  of  oxygen,  and  0.94  per  cent  of  argon.  When  only  the 
water  is  removed,  the  carbon  dioxide  averages  about  0.03  per  cent 
of  the  whole. 

To  use  an  illustration  of  Graham's,  if  we  imagined  the  air  to  be 
divided  by  magic  into  layers,  all  at  one  atmosphere  pressure,  and 
with  the  heavier  components  below,  we  should  have:  On  the  earth, 
five  inches  of  water;  above  that,  thirteen  feet  of  carbon  dioxide; 
above  that,  ninety  yards  of  argon;  above  that,  one  mile  of  oxygen; 
and  on  the  top  four  miles  of  nitrogen. 

Liquefaction  of  Gases.  —  The  earliest  experiments  of  this  kind 
seem  to  have  been  made  by  Northmore  (1805),  who  liquefied  chlorine, 
hydrogen  chloride,  and  sulphur  dioxide.  In  1823  chlorine  was  again 
liquefied  by  Faraday;  and  in  the  same  year  Davy,  whose  assistant 
Faraday  was,  liquefied  hydrogen  chloride.  During  the  following  years 
Faraday  reduced  other  gases  —  sulphur  dioxide,  hydrogen  sulphide, 
carbon  dioxide,  nitrous  oxide,  cyanogen,  and  ammonia  —  to  the  liquid 
condition.  He  failed,  however,  with  oxygen,  hydrogen,  and  nitrogen. 

The  method  which  he  employed  was  extremely  simple.  He  used  a 
bent  tube  shaped  like  an  inverted  v  (A),  into  one  limb  of  which 
materials  for  producing  the  gas  were  introduced  (Fig.  90,  p.  226). 
The  other  limb  was  then  sealed  up  and  immersed  in  a  freezing  mix- 
ture. The  gas,  usually  liberated  by  heating,  was  liquefied  by  its  own 
pressure  in  the  cold  limb.  By  means  of  a  more  elaborate  apparatus, 
Cailletet  and  Pictet  simultaneously  (December,  1877)  obtained,  the 
one,  a  fog,  and  the  other,  a  spray  containing  droplets  of  liquid  oxygen. 
In  1883  Wroblevski  and  Olszevski  made  visible  amounts  of  the  same 
liquid.  About  the  same  time  Dewar  devised  means  of  manufacturing 
large  quantities  of  liquid  air  and  oxygen. 

The  principle  now  used  in  liquefying  gases  depends  on  the  fact 
that,  although  a  perfect  gas,  when  expanding  into  a  vacuum,  should 
suffer  no  fall  in  temperature,  since  it  does  no  work,  ordinary  gases  do 
become  cooled  very  slightly.  The  work  which  they  do  in  expand- 
ing in  such  circumstances  is  done  in  overcoming  the  cohesion  between 
their  molecules  (p.  164),  so  that  a  tearing  apart  of  the  substance, 
which  consumes  heat,  has  to  take  place.  Since  this  cohesion  becomes 
more  conspicuous  the  lower  the  temperature  (cf.  p.  166),  the  cooling 
effect  of  expansion  becomes  greater  and  greater  as  the  temperature 
falls. 


508 


INORGANIC   CHEMISTRY 


The  most  successful  apparatus  for  use  on  a  small  or  large  scale  is 
that  devised  by  Hampson.  In  this  apparatus  (Fig.  117),  two  con- 
centric copper  pipes,  about  130  meters  in  length,  are  coiled  closely  in 
a  cylindrical  form,  with  non-conducting  covering  to  prevent  access  of 
heat  from  the  outside.  Air  at  130-150  atmospheres  pressure  is 
forced  through  the  inner  pipe  (upper  opening,  Fig.  117).  When  it 
reaches  the  extremity  of  this  pipe,  it  suddenly  escapes  into  a  closed 
vessel.  This  expansion  lowers  its  temperature.  A  spiral  partition 


FIG.   117. 


FIG.   118. 


between  the  coils  produces  the  outer  tube  of  which  we  have  spoken. 
The  gas  in  the  tube  A  (Fig.  118)  is  under  a  pressure  of  130-150 
atmospheres.  The  distance  of  the  nozzle  D  from  the  plug  C  is 
adjusted  so  that  the  pressure  of  the  gas  in  the  chamber  and  spiral 
outer  tube  is  reduced  to  one  atmosphere.  The  air  can  now  escape 
only  by  traveling  back  through  the  outer  pipe  to  the  final,  wider  exit 
near  the  top.  In  doing  so,  it  cools  the  highly  compressed  air  in  the 
inner  pipe.  The  cooler  air,  on  reaching  the  closed  vessel,  expands 
and  becomes  colder  than  ever,  and  in  passing  backwards  lowers  the 
temperature  of  the  air  in  the  inner  pipe  still  further.  Finally,  the  air 
in  this  pipe  liquefies,  and  drops  of  liquid  air  are  expelled  into  the 


THE  ATMOSPHERE.      THE  HELIUM   FAMILY  509 

closed  vessel.     This  is  allowed  to  run  out  through  a  valve,  from 
time  to  time,  as  it  accumulates. 

Liquid  air  can  be  kept  in  Dewar  flasks  (Fig.  119).  The  space 
between  the  inner  and  outer  flasks  is  evacuated,  so  that  there  is  no 
gas  to  carry  heat  from  the  atmosphere  in  to  the  liquid  air.  The  inner 
surface  of  the  outer  flask  is  often  silvered,  so  that  radiant 
heat,  from  surrounding  bodies,  may  be  reflected  and  not 
absorbed. 


Liquid  Air.  —  Liquid  air  varies  in  composition,  as 
the  nitrogen  (b.-p.  —194°)  is  less  condensible  than  the 
oxygen  (b.-p.  -181.4°).  It  boils  at  about  -190°,  and 
contains  about  54  per  cent  of  oxygen  by  weight,  while  air  contains 
23.2  per  cent.  By  allowing  evaporation  to  go  on,  a  liquid  containing 
75  to  95  per  cent  of  oxygen  is  easily  obtained  (cf.  p.  81).  The  gas 
secured  by  the  evaporation  of  the  residue  is  pumped  into  cylinders 
and  sold  as  compressed  oxygen.  It  contains  about  3  per  cent  of 
argon,  and  is  a  convenient  source  of  this  element.  Cartridges  made 
of  granular  charcoal  and  cotton  waste,  when  saturated  with  liquid 
air,  have  been  used  as  an  explosive  in  mining. 

THE  HELIUM  FAMILY 

Argon  A.  —  Lord  Rayleigh  was  the  first  to  observe  that,  while 
specimens  of  oxygen  and  other  gases  made  purposely  from  various 
sources  always  had  the  same  density,  nitrogen  was  an  exception. 
One  liter  of  nitrogen  made  from  air,  and  supposed  to  be  pure,  weighed 
1.2572  g.  When  the  gas  was  manufactured  by  decomposition  of  five 
different  compounds,  such  as  urea  and  certain  oxides  of  nitrogen,  the 
results  agreed  well  amongst  themselves.  But  the  mean  weight  of  a 
liter  of  this  nitrogen  was  only  1.2505  g.  The  difference,  amounting  to 
nearly  7  mg.,  was  very  much  greater  than  the  experimental  error. 
The  suspicion  naturally  arose  that  some  heavier  gas  was  present  in 
atmospheric  nitrogen.  Soon  after  (1894),  the  late  Sir  William  Ram- 
say, working  in  cooperation  with  Lord  Rayleigh,  obtained  argon  by 
removal  of  the  greatly  preponderating  nitrogen  by  means  of  mag- 
nesium (p.  514).  The  new  gas  had  a  molecular  weight  of  about  40, 
and  was  therefore  more  than  one-third  heavier  than  nitrogen. 

In  order  to  make  sure  that  this  substance  did  not  have  its  source  in  the  magne- 
sium, a  different  method  was  used  by  Lord  Rayleigh  to  separate  it  from  nitrogen. 


510  INORGANIC  CHEMISTRY 

He  inclosed  the  nitrogen  with  a  sufficient  quantity  of  oxygen  in  a  flask,  through 
the  sides  of  which  platinum  poles  had  been  inserted.  A  tube  entered  the  flask  by 
the  neck,  and  through  this  a  constant  fountain  of  potassium  hydroxide  solution 
played  upon  the  interior  and  kept  the  surface  covered  with  fresh  quantities  of  the 
liquid.  Another  tube  permitted  the  overflow  of  the  excess  of  this  solution.  The 
discharge  of  electricity  produced  nitrogen  tetroxide  (q.v.),  which  was  absorbed  by 
the  potassium  hydroxide  to  form  potassium  nitrate  and  potassium  nitrite.  The 
volume  of  gas  thus  continually  diminished  and,  by  persistent  sparking  of  the  mix- 
ture with  oxygen,  the  nitrogen  was  finally  all  taken  out.  The  excess  of  oxygen 
was  then  removed,  and  the  gas  which  remained  was  found  to  be  identical  with  that 
which  Ramsay  had  obtained. 

Lord  Rayleigh's  method  was  extremely  interesting,  since  it  was  a  reproduction 
of  an  experiment  made  by  Cavendish  more  than  one  hundred  years  before  (1785). 
The  latter  had  remarked  that  the  assumption  that  the  inert  atmospheric  gas  was 
a  homogeneous  single  substance  had  not  been  confirmed  by  sufficiently  careful 
experiment.  He  even  endeavored  in  precisely  the  above  way  to  remove  the  nitro- 
gen in  order  to  see  whether  any  other  body  remained.  He  records  the  fact  that 
only  about  0.8  per  cent  of  inactive  gas  remained.  Since  the  quantity  was  so 
small,  and  the  spectroscope,  by  which  the  gas  even  in  small  amounts  would  have 
been  recognized  to  be  new,  was  not  invented  until  much  later,  he  did  not  pursue 
the  subject.  Argon  thus  narrowly  escaped  detection  over  a  century  before  its 
actual  discovery. 

The  exact  density  of  argon,  referred  to  oxygen  =  32,  is  39.88. 
When  liquefied  it  boils  at  — 186.9°,  and  the  colorless  solid,  obtained 
by  cooling  the  liquid,  melts  at  — 189.5°.  The  solubility  of  the  gas  in 
water  (4  volumes  in  100)  is  two  and  one-half  times  that  of  nitrogen. 
It  has  not  been  found  to  enter  into  any  sort  of  chemical  combination, 
and  was  named  argon  on  this  account  (Gk.  apyos,  inactive). 

Since  the  atomic  weight  of  a  substance  is  a  quantity  showing  the 
proportion  in  which  it  enters  into  combination,  it  will  be  seen  that 
argon,  since  it  has  not  yet  been  found  to  combine  with  anything,  has, 
to  speak  strictly,  no  atomic  weight  (pp.  243,  254).  At  the  same  time, 
it  is  manifestly  a  question  of  interest  to  determine  whether  the  physi- 
cal properties  require  the  supposition  that  the  molecule  of  argon 
contains  one  atom,  or  more  than  one  atom. 

If  gases  were  composed  of  perfectly  elastic  spheres,  the  molecules 
would  be  altered  only  in  respect  to  velocity  of  movement  by  heating. 
Calculation  enables  us  to  determine  that  to  raise  the  temperature  of 
one  G.M.V.  of  such  a  gas  by  one  degree  would  require  3  calories  in 
every  case.  Now  Regnault  found  the  following  values  (in  calories) 
for  the  heat  capacity  of  gases: 


THE  ATMOSPHERE.     THE  HELIUM  FAMILY  511 

Oxygen  (O2) 4.96  Carbon  dioxide  (C02)     .     7.56 

Hydrogen  (H2) 4.82  Sulphur  dioxide  (S02)      .     7.82 

Nitrogen  (N2)      4.82  Chloroform  (CHC13)    .    .   16.55 

Hydrogen  chloride  (HC1)  .  4.76  Alcohol  (C2H60)  ....  18.70 

These  gases  evidently  are  not  constituted  as  the  hypothesis  with 
which  this  paragraph  opened  supposes.  Some  heat  is  consumed  in 
work  done  inside  the  polyatomic  molecules,  and  the  amounts  by  which 
the  numbers  exceed  3  calories  show  that  the  intramolecular  work  is 
greater  as  the  complexity  of  the  molecules  increases.  Now,  in  mer- 
cury vapor  the  value  is  exactly  3,  and  its  vapor  is  monatomic  (p.  250), 
so  that  in  its  molecules  there  is  no  opportunity  for  the  consumption 
of  heat  in  intramolecular  change.  Hence,  when  argon  was  found 
likewise  to  give  3  for  the  value  of  its  molecular  heat-capacity,  identity 
of  its  atomic  and  molecular  weights  was  demonstrated. 

Helium  He.  —  In  1868  Lockyer  first  detected  an  orange  line  in 
the  spectrum  of  the  sun's  prominences,  which  was  not  given  by  any 
terrestrial  substance  then  known.  The  line  was  so  conspicuous  that 
it  was  attributed  to  the  presence,  in  considerable  quantity,  of  a  new 
chemical  element,  which  was  named  helium  (Gk.  rj\ios,  the  sun). 
Ramsay,  in  searching  for  sources  of  argon,  examined  a  gas  which 
Hillebrand  had  obtained  by  heating  uraninite,  an  ore  of  uranium.  He 
was  surprised  to  find  (1895)  that  the  gas  was  not  nitrogen,  as  had  been 
supposed,  nor  was  it  even  argon.  It  frequently  contained  a  large 
proportion  of  a  gas,  very  much  lighter  than  either,  the  spectrum  of 
which  showed  at  once  that  it  was  identical  with  helium.  The  same 
gas  has  since  been  obtained  from  other  minerals,  from  the  water  of 
certain  mineral  springs,  and  it  is  found  in  small  amount  in  the  atmos- 
phere. Helium  does  not  exhibit  any  tendency  to  enter  into  com- 
bination, either  with  the  elements  which  its  parent  minerals  contain, 
or  with  any  others.  It  is  monatomic  (cf.  p.  510)  and  its  density 
shows  that  its  molecular  weight  is  4.  When  liquefied  by  Onnes,  it 
boiled  at  -268.5°  (4.5°  abs.),  and  had  a  density  of  only  0.15. 

Neon  Ne,  Krypton  Kr,  Xenon  Xe,  and  Niton  Nt.  —  When 
the  argon  obtained  from  atmospheric  nitrogen  was  cooled  with  liquid 
air  (—190°)  by  Ramsay  (1898),  the  argon,  krypton,  and  xenon  were 
liquefied,  and  the  neon  and  helium  were  dissolved  by  the  liquid. 
When  heat  was  allowed  to  reach  the  mixture,  the  last  two  gases 
escaped  first,  along  with  much  argon.  When  most  of  the  argon  had 


512  INORGANIC   CHEMISTRY 

escaped,  the  krypton  and  xenon  still  remained  liquid.  By  repeated 
liquefaction  and  fractional  evaporation  (see  under  Petroleum),  the 
krypton  and  xenon  were  separated  from  the  argon  and  from  one 
another.  When  the  vessel  containing  the  mixture  of  helium  and  neon 
was  immersed  in  liquid  hydrogen  (  —  240°),  the  second  froze  to  a  white 
solid,  and  the  helium,  which  remained  gaseous,  could  be  pumped  off. 
All  these  gases  together  constituted  one-six  hundredth  part  of  the 
atmospheric  argon. 

These  gases  are  all  entirely  inactive  chemically,  and  are  all  mon- 
atomic.  Their  molecular  weights  are:  Neon  (Gk.  vfa,  new),  20.2; 
krypton  (Gk.  Kpvirrov,  hidden),  82.92;  xenon  (Gk.  £evov,  stranger), 
130.2. 

Niton  (radium  emanation,  q.v.),  molecular  weight  222.4,  also 
belongs  to  this  family. 

Exercises.  —  1.  A  sample  of  moist  air,  confined  over  water  at 
15°  and  760  mm.,  occupies  15  c.c.  It  is  mixed  with  20  c.c.  of  hydro- 
gen, and  the  mixture  is  exploded,  and  suffers  a  contraction  of  9.5  c.c. 
What  would  be  the  volume  of  the  oxygen  it  contained  if  measured  dry 
at  0°  and  760  mm.? 

2.  Calculate,  from  the  data  on  p.  500  and  the  densities,  the 
percentage  by  weight  of  the  three  principal  components  of  air. 

3.  Of  the  proofs  that  air  is  a  mixture  (p.  506),  which  show  that 
no  part  of  the  components  is  combined,  and  which  that  the  com- 
ponents are  not  wholly  combined? 

4.  What  is  the  relation  between  heavier  clothing  and  the  station- 
ary layer  of  air  next  the  skin? 

5.  Prom  the  data  given  on  p.  502,  calculate  the  weight  of  water 
vapor  in  1  cubic  meter  of  air  saturated  at  18°  and  at  0°,  respectively . 


CHAPTER  XXV 
NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN 

NITROGEN  was  recognized  to  be  a  distinct  substance  by  Ruther- 
ford (1772),  Professor  of  Botany  in  the  University  of  Edinburgh, 
who  named  it  mephitic  air.  Scheele  showed  that  it  was  present 
in  the  atmosphere.  Lavoisier  recognized  it  to  be  an  element,  and 
named  it  azote  (Gk.,  without  life)  because  it  did  not  support  life. 
The  English  name  records  the  fact  that  it  is  an  important  con- 
stituent of  saltpeter  KN03  (Lat.  nitrum). 

The   Chemical  Relations  of  the  Element  Nitrogen.  —  In 

compounds  with  hydrogen  and  the  metals  nitrogen  is  trivalent, 
while  in  those  containing  oxygen  and  other  negative  elements  it  is 
frequently  quinquivalent.  It  is  a  non-metal,  for  its  oxides  are 
acidic  (p.  150).  Many  of  the  compounds  of  nitrogen  are  extremely 
active  and  interesting.  Those  of  them  which  we  have  to  discuss 
in  inorganic  chemistry  are  ammonia  NH3  and  nitric  acid  HNO3,  and 
several  related  substances.  The  organic  compounds  containing 
nitrogen  are  very  numerous  and  possess  highly  characteristic  prop- 
erties. Some,  like  nitroglycerine  (q.v.)  and  guncotton,  are  violently 
explosive;  others,  like  antipyrine,  show  great  physiological  activity; 
still  others,  such  as  the  aniline  and  other  organic  dyes,  provide  us  with 

beautiful  and  useful  coloring  matters. 

. 

Occurrence.  —  Apart  from  the  presence  of  free  nitrogen  in  the 
air,  the  element  is  found  in  many  forms  of  combination.  The  nitrates 
of  potassium  and  sodium  are  found  in  Bengal  and  other  tropical 
countries  and  in  Chile,  respectively.  Natural  manures,  such  as 
guano,  contain  large  quantities  of  nitrogen  compounds,  and  owe  their 
value  as  fertilizers  to  this  fact.  Nitrogen  is  an  essential  constituent 
of  the  proteins  (about  16  per  cent  nitrogen)  of  vegetable  and  animal 
matter. 

Preparation.  —  Nitrogen  containing  about  one  per  cent  of  argon 
(q.v.)  is  easily  obtainable  from  purified  air,  when  the  oxygen  of  the 

513 


514  INORGANIC   CHEMISTRY 

latter  is  removed,  for  example,  by  passage  over  heated  copper.  For 
commercial  purposes,  it  is  obtained  by  evaporation  of  liquid  air. 

When  pure  nitrogen  is  required,  it  must  be  obtained  from  chemical 
compounds,  in  order  that  it  may  be  free  from  argon.  The  simplest 
method  is  to  heat  ammonium  nitrite: 

NH4N02->2H20  +  N2. 

In  practice,  since  ammonium  nitrite  is  unstable  and  cannot  easily  be 
kept,  a  mixture  of  an  ammonium  salt  with  some  salt  of  nitrous  acid  is 
employed.  Thus,  when  strong  solutions  of  ammonium  chloride  and 
sodium  nitrite  are  mixed,  a  double  decomposition  results  in  the  for- 
mation of  ammonium  nitrite,  NH4C1  +  NaNO2  +±  NH4NO2  +  NaCl, 
and  this  breaks  up  when  heat  is  applied,  giving  nitrogen. 

We  may  also  prepare  nitrogen  by  the  oxidation  of  ammonia  NH3, 
by  passing  the  latter  over  heated  cupric  oxide  (see  p.  519),  or  by  the 
reduction  of  nitric  oxide  NO  by  passing  this  gas  over  heated  copper. 

Physical  Properties.  —  Nitrogen  is  a  colorless,  tasteless,  odor- 
less gas,  as  we  should  expect  from  the  fact  that  air  possesses  these 
properties.  It  forms  a  colorless  liquid,  boiling  at  —194°.  By 
further  cooling,  this  liquid  freezes  to  a  white  solid  (m.-p.  —214°). 
The  solubility  of  nitrogen  in  water  (1.6  vols.  in  100)  is  less  than  that 
of  oxygen. 

Chemical  Properties.  —  The  density  of  the  gas  shows  the 
formula  of  free  nitrogen  to  be  N2. 

Nitrogen  unites  with  few  common  chemical  elements  directly.  At 
ordinary  temperatures  it  is  almost  absolutely  indifferent.  When 
passed  through  a  tube  over  strongly  heated  lithium,  calcium,  magne- 
sium, or  boron,  it  forms  definite  chemical  compounds,  known  as 
nitrides,  in  which  it  is  trivalent  These  have  the  formulae  Li3N,  Ca3N2, 
Mg3N2,  and  BN,  respectively.  Thus,  when  magnesium  is  burned  in 
the  air,  the  white  mass  which  is  formed  contains  magnesium  nitride, 
along  with  much  of  the  oxide.  When  the  ash  is  moistened  with  water 
in  a  covered  vessel,  ammonia  can  be  smelt  and  can  be  detected  with 
moist  litmus  paper  [Lect.  exp.].  The  nitride  is  hydrolyzed: 

Mg3N2  +  6H20  ->  3Mg(OH)2  +  2NH3  T  . 

Nitrogen  combines  with  difficulty  with  hydrogen  to  form  ammonia 
NH3  and  with  still  greater  difficulty  with  oxygen  to  form  nitric  oxide 
NO.  The  actions  will  be  discussed  under  the  compounds  themselves. 


NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN     515 

The  indifference  of  free  nitrogen  is  doubtless  due  to  the  fact  that  its 
molecules  N2  are  extremely  stable. 

One  case  of  direct  union  of  nitrogen  is  of  economic  importance. 
The  supply  required  by  plants  is  obtained  partly  from  nitrogen  com- 
pounds contained  in  fertilizers,  or  equivalent  substances  already  pres- 
ent in  the  soil.  The  leguminosce,  such  as  peas,  beans,  clover,  and 
alfalfa,  are  associated  constantly  with  certain  bacteria  which  flourish 
in  nodules  upon  their  roots.  These  bacteria  have  the  power  of  taking 
free  nitrogen  from  the  air,  which  penetrates  the  soil,  breaking  up  its 
molecules,  and  producing  proteins.  The  nodules  often  contain  over 
five  per  cent  of  combined  nitrogen.  The  proteins,  by  the  action  of 
nitrifying  bacteria,  give  nitric  acid  which,  with  bases  in  the  soil, 
gives  nitrates.  These  are  soluble,  and  are  absorbed  through  the 
roots,  furnishing  the  nitrogen  needed  by  plants  to  enable  them  to 
construct  the  proteins  they  require. 

An  active  form  of  nitrogen,  discovered  by  Strutt,  is  produced  by 
passing  an  electric  discharge  through  the  gas  under  low  pressure. 
When  the  discharge  is  stopped,  a  yellow  light  is  produced  by  the 
reversion  of  the  nitrogen  to  the  inactive  form.  A  trace  of  oxygen 
seems  to  be  required,  possibly  as  a  catalyst.  With  a  little  vapor 
of  pentane  CsH^  in  the  tube,  this  active  nitrogen  gives  hydrocyanic 
acid  HNC  (q.v.). 

Compounds  of  Nitrogen  and  Hydrogen.  —  The  commonest 
and  longest  known  of  these  substances  is  ammonia  NH3,  which  was 
first  described  by  Priestley  (1774)  and  named  "alkaline  air. "  Curtius 
discovered  hydrazine  N2H4  in  1889,  and  hydrazoic  acid  HN3  in  1890. 
Hydroxylamine  NH3O,  discovered  by  Lossen  in  1865,  is  similar  to 
ammonia  in  chemical  behavior. 

AMMONIA  NH3 

Ammonia  is  of  interest,  commercially,  because  large  amounts 
of  liquefied  ammonia  are  used  in  refrigeration,  because  much  is 
employed  in  the  manufacture  of  carbonate  of  soda,  and  because  its 
compounds  are  used  as  fertilizers. 

Manufacture.  —  Ammonia  is  formed  when  proteins  are  heated 
in  the  absence  of  air.  Thus,  it  was  formerly  obtained  by  the  distilla- 
tion of  hoofs,  hides,  and  horns,  and  the  solution  in  water  was  called 
" spirit  of  hartshorn."  Coal  contains  1-2  per  cent  of  combined 


516  INORGANIC  CHEMISTRY 

nitrogen,  derived  from  the  proteins  of  the  original  plants.  Hence, 
when  coal  is  distilled  in  the  manufacture  of  coal  gas  or,  on  a  much 
larger  scale,  for  the  making  of  coke,  much  ammonia  can  be  secured 
by  washing  with  water  the  gases  which  are  given  off.  The  solution 
is  separated  from  the  tar,  lime  is  added  to  liberate  the  ammonia,  and 
the  ammonia  gas  is  driven  out  by  heating  and  passed  into  sulphuric 
acid  or  hydrochloric  acid.  It  gives  ammonium  sulphate  or  chloride 
(see  below). 

In  Germany  80  per  cent  (1910)  of  the  coke  is  made  in  "by- 
product" coke  ovens,  in  which  the  ammonia  and  other  by-products 
are  collected  and  utilized;  in  the  United  States  83  per  cent  of  the 
coke  is  made  in  "beehive"  ovens  (1911)  in  which  the  vapors  are 
simply  burned.  Ammonium  sulphate  is  a  valuable  fertilizer  and  in 
1911,  in  the  United  States,  ammonia  capable  of  yielding  400,000 
tons  of  ammonium  sulphate  worth  24  million  dollars  was  burned 
by  the  cokemakers. 

The  distillation  of  coal  is  the  chief  source  of  commercial  am- 
monia. In  Scotland,  however,  oil-bearing  shale  is  distilled  to  obtain 
petroleum,  and  much  ammonia,  liberated  at  the  same  time,  is  col- 
lected. Formerly  it  was  allowed  to  escape  but,  in  the  absence  of  a 
protective  tariff,  the  competition  of  American  and  Russian  petroleum 
compelled  economy.  Now,  the  profit  on  the  ammonium  sulphate 
pays  the  whole  cost  of  mining  and  distilling  the  shale. 

Synthetic  Ammonia.  —  The  direct  union  of  nitrogen  and 
hydrogen  (1  vol.:  3  vols.)  to  form  ammonia:  N2  +  3H2  <=±  2NH3  + 
2  X  12,200  cal.  is  a  reversible  reaction.  Since  heat  is  absorbed  in 
decomposing  ammonia,  the  proportion  of  this  gas  present  at  equi- 
librium becomes  rapidly  smaller  as  the  temperature  rises  (van't 
Hoff's  law,  p.  305).  The  proportions  (at  760  mm.),  determined  by 
Haber,  are:  200°,  15.3  per  cent;  300°,  2.2  per  cent;  500°,  0.13  per 
cent;  700°,  0.02  per  cent;  1000°,  0.004  per  cent.  Thus,  at  700° 
ammonia  is  practically  all  decomposed,  and  at  low  temperatures  the 
action  is  so  slow  that  no  union  to  form  ammonia  is  perceptible. 

The  Badische  Company  is  now  (1914)  manufacturing  ammonia 
on  a  large  scale,  for  the  preparation  of  explosives,  by  the  direct  union 
of  nitrogen  and  hydrogen.  It  is  necessary  to  use  a  lower  temperature 
and  a  contact  agent  —  such  as  specially  prepared  iron  —  to  hasten 
the  action.  Then,  too,  the  reaction  is  accompanied  by  a  diminution 
in  volume  (4  vols.  — » 2  vols.),  and  is  therefore  assisted  by  using  the 
under  a  pressure  of  185-200  atmospheres  (Le  Chatelier's  law, 


NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN     517 

p.  307).  Below  500°,  with  these  conditions,  about  8  per  cent  of  the 
gases  combine.  The  ammonia  is  dissolved  out  with  water,  and 
the  uncombined  gases  are  sent  through  the  process  again.  The 
required  hydrogen  may  be  obtained  by  one  of  the  commercial  proc- 
esses (p.  122),  and  the  nitrogen  from  liquid  air. 

Preparation  in  the  Laboratory.  —  1.  A  mixture  of  slaked 
lime  and  some  salt  of  ammonium,  such  as  ammonium  chloride, 
either  with  or  without  water,  is  heated  in  a  flask  or  retort  provided 
with  a  delivery  tube : 

Ca(OH)2  +  2NH4C1  *=?  CaCl2  +  2NH4OH  «=>  2NH3  +  2H2O. 

The  ammonium  hydroxide,  formed  by  the  double  decomposition, 
being  unstable,  immediately  decomposes. 

2.  Warming  the  aqueous  solution  gives  a  steady  stream  of  the 
gas.  Since  the  gas  is  very  soluble  in  water,  it  is  collected  over 
mercury  or  by  downward  displacement  of  air.  The  moisture  which 
it  contains  is  removed  by  passage  through  a  tower  or  wide  tube 
filled  with  quicklime.  Calcium  chloride  cannot  be  used,  as  the  gas 
combines  with  it  forming  a  compound  CaCl2,8NH3,  similar  in  prop- 
erties to  a  hydrate  (pp.  150-154). 

In  the  decay  of  meat  (containing  proteins)  and  of  manure,  the 
strong  odor  is  due  in  part  to  the  ammonia  produced. 

Physical  Properties.  —  Ammonia  is  a  colorless  gas  with  a  pun- 
gent, characteristic  odor  familiar  in  smelling-salts.  The  G.M.V.  of 
the  gas  weighs  17.26  g.,  so  that  the  density  is  little  more  than  half 
that  of  air  (cf.  p.  233) .  When  liquefied  it  boils  at  —  33°  and  exercises 
a  pressure  of  6.5  atmospheres  at  10°  (density  0.62  at  0°).  The  solid 
is  white  and  crystalline  (m.-p.  —77°).  The  gas  is  very  soluble  in 
water,  one  volume  of  the  latter  dissolving  1300  volumes  of  the  gas  at 
0°,  783  volumes  at  16°,  and  306  volumes  at  50°.  The  35  per  cent 
aqueous  solution,  saturated  at  15°,  and  sold  as  "  concentrated 
ammoni^, "  has  a  sp.  gr.  0.881.  The  whole  of  the  dissolved  gas  may 
be  removed  by  boiling  (cf.  p.  211). 

Liquefied  ammonia  is  used  in  refrigeration.  In  evaporating  at 
—  33°  it  absorbs  330  cal.  per  gram.  Water  alone  has  a  greater 
heat  of  vaporization.  The  large  amount  of  heat  is,  in  both  cases, 
required  because  of  the  relatively  large  volume  of  the  vapor  (due  to 
low  molecular  weight)  and  to  the  fact  that  both  liquids  are  associated 
(p.  282),  and  the  complex  molecules  (NH3)2  and  (NH3)3  have  to  be 


518 


INORGANIC  CHEMISTRY 


decomposed.     To  freeze  1  gram  of  water  at  0°,  80  cal.  have  to  be 
removed.     Thus  1  g.  of  liquid  ammonia  will  convert  4  g.  of  water 

„ into  ice.     Fig.  120  shows  one  arrangement 

diagrammatically.  The  ammonia  gas,  ob- 
tained from  a  cylinder  of  liquid  ammonia, 
is  driven  by  the  pump  F  along  the  tube  E 
and  is  liquefied  in  the  tube  coiled  in  the 
tank  AB.  Cold  water  circulating  through 
AB  removes  the  heat  produced  by  the  com- 
pression and  liquefaction  of  the  gas.  The 
liquid  ammonia  is  allowed  to  drip  through 
the  stopcock  G  into  the  lower  coil,  and  there 
it  evaporates.  In  doing  so,  it  takes  heat 
from  a  30  per  cent  solution  of  calcium  chlo- 
ride in  water.  This  cooled  brine  leaves  the 
tank  at  D,  circulates  through  another  tank, 
in  which  water-filled  ice  molds  are  sus- 
pended, and  returns  to  C.  When  used  for 
cooling  storage-rooms  for  meat,  the  brine 
circulates  through  pipes  in  the  same  way. 
The  machine  is  constructed  of  iron,  because 


FIG.  120. 


copper  and  brass  are  corroded  by  ammonia. 

Chemical  Properties.  —  Ammonia,  as  we  have  seen 
(p.  516),  is  not  stable,  and  decomposes  almost  completely 
at  700°.  A  discharge  of  sparks  from  an  induction  coil 
(temperature  about  2000°)  has  the  same  effect,  so  that 
a  sample  of  the  gas,  confined  over  mercury  in  a  closed 
tube  (Fig.  121),  may  be  shown  to  double  in  volume. 
Every  two  molecules  give  four: 

2NH3  <=±  3H2  +  N2. 

That,  even  at  this  temperature,  the  action,  being  re- 
versible, is  still  incomplete,  can  be  shown  by  introduc- 
ing a  few  drops  of  dilute  sulphuric  acid.  The  trace  of 
ammonia  remaining  combines  with  this  acid,  forming 
(NH4)2S04  in  solution.  If  the  discharge  is  continued, 
further  traces  of  ammonia  are  formed  and  absorbed,  until,  finally, 
the  whole  gas  disappears.  The  action,  therefore,  first  goes  almost 
completely  in  one  direction,  and  then  quite  completely  in  the  other, 
while  no  change  has  taken  place  in  the  conditions  to  which  the  gas 


FIG.  121. 


NIlftOGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN     519 

is  subjected  at  the  point  where  the  interaction  is  occurring.  The  sole 
difference  is  that  a  little  of  an  acid  has  been  introduced  into  a  rela- 
tively remote  part  of  the  space. 

Ammonia  reduces  many  oxides,  when  the  latter  are  heated  and 
the  gas  is  led  over  them: 

3CuO  +  2NH3  ->  3Cu  +  3H2O  +  N2. 

Ammonia  burns  in  pure  oxygen  (not  in  air)  to  give  steam  and  nitrogen. 
In  air,  heat  is  used  up,  not  only  in  decomposing  the  compound,  but 
also  in  raising  the  temperature  of  the  nitrogen  of  the  atmosphere. 
This  causes  a  continuous  drain  on  the  heat  given  out  by  the  action, 
and  prevents  the  maintenance  of  the  kindling  conditions  (p.  95). 

Chlorine  and  bromine  (vapor)  combine  with  the  hydrogen  and 
liberate  nitrogen: 

2NH3  +  3C12  -*  N2  +  6HC1. 

When  metals  capable  of  uniting  with  nitrogen  (p.  514)  are  heated 
in  a  stream  of  ammonia  gas,  hydrogen  is  displaced.  Magnesium 
gives  magnesium  nitride: 

2NH3  +  3Mg  ->  Mg3N2  +  3H2. 

Sodium  and  potassium,  however,  give  amides  (compounds  con- 
taining the  group  NH2),  such  as  sodamide  NaNH2: 

2NH3  +  2Na  -»  2NaNH2  +  H2. 

The  most  striking  property  of  ammonia  is  that  it  combines  with 
acids,  giving  ammonium  salts : 

NH3  (gas)  +  HC1  (gas)     -»  NH4C1  (solid). 
2NH3  (gas)  +  H2S04  (liq.)  -»  (NH4)2S04  (solid). 

It  combines  also  with  water  at  or  below  —79.3°  to  give  ammonium 
hydroxide  as  a  white  solid: 

NH3  +  H20  ->  NH4OH. 

Since  the  solid  liquefies  above  —79.3°,  a  solution  of  the  substance, 
which  is  contained  in  the  aqueous  solution  of  ammonia,  is  the  only 
available  form  of  ammonium  hydroxide.  In  solution,  it  is  a  weak 
base.  Ammonium  oxide  (NH4)20,  a  solid,  can  also  be  formed  below 
-78.6°. 

Ammonium  Compounds.  —  Since  NH4  plays  the  part  of  a 
metallic  element,  entering  into  the  composition  of  a  base  and  of 


520  INORGANIC  CHEMISTRY 

a  series  of  salts,  it  is  named  ammonium.  As  this  radical  forms  a 
univalent,  positive  ion  NH4+  and  gives  a  distinctly  alkaline  base,  it 
is  classed  with  potassium  and  sodium  as  one  of  the  metallic  elements 
of  the  alkalies  (q.v.). 

Ammonium  Hydroxide  NH^OH.  —  Although  much  less  com- 
pletely ionized  than  is  potassium  hydroxide,  ammonium  hydroxide 
affects  litmus  easily.  In  a  normal  solution  0.4  per  cent  of  the 
ammonia  is  in  the  form  of  ammonium-ion  NH4+.  When  an  acid  is 
added  to  the  solution,  the  correspondingly  small  amount  of  hydrox- 
ide-ion which  exists  in  it  is  removed,  and  the  various  equilibria  are 
displaced  forwards.  The  final  result  is  the  same  as  with  any  other 
base: 


NH3  (gas)^±NH3  (dslvd) 

- 


Only  a  small  proportion  of  the  gas  (one-third)  is  actually  com- 
bined at  any  one  time,  the  greater  part  being  simply  dissolved. 

The  solution  is  sold  as  household  ammonia,  and  is  used,  in  wash- 
ing and  cleaning,  to  soften  the  water. 

Salts  of  Ammonium.  —  When  strongly  heated,  all  ammonium 
salts  are  decomposed  and  many,  but  not  all,  give  ammonia  and  the 
acid.  When  the  latter  is  volatile,  the  whole  material  of  the  salt  is 
thus  converted  into  gas.  If  the  acid  is  volatile  without  permanent 
decomposition,  it  reunites  with  the  ammonia  to  form  the  solid  salt 
when  the  vapor  reaches  a  cool  part  of  the  tube  (sublimation,  p.  275)  : 

NH4C1  (solid)  <=»  NH4C1  (gas)  <=±  HC1  +  NH3. 

This  behavior  distinguishes  ammonium  salts  from  those  of  the  typical 
metals,  for,  with  the  exception  of  mercury  salts,  most  other  salts 
are  not  easily  and  completely  volatilized.  The  use  of  ammonium 
chloride  (sal  ammoniac)  in  soldering  depends  on  the  dissociation  of 
the  salt,  by  the  heat  of  the  iron,  and  the  action  of  the  liberated  hydro- 
gen chloride  on  the  oxide  which  covers  the  surface  of  the  metal  to 
be  soldered. 

Some  ammonium  salts,  like  the  nitrite  (p.  514),  when  heated, 
give  no  ammonia  (see  also  nitrous  oxide  and  ammonium  dichromate). 
Conversely,  substances  containing  proteins  (e.g.,  gelatin),  when 
heated,  do  give  ammonia.  Hence,  liberation  of  ammonia  is  not 
proof  that  a  substance  is  a  salt  of  ammonium. 


NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN     521 

The  test  for  ammonium  salts  is  to  warm  them,  dry  or  in  solution, 
with  a  base,  when  the  odor  of  ammonia  becomes  noticeable: 

(NH4)2S04 1*  S04=  +  2NH4+  U 

i    2OH~  (  ±14U±1  ?=±  J±12U  -f-  ^JN  rl3  |  . 


When  the  solution  is  used,  it  is  the  tendency  of  the  NH4+  and  OH~ 
to  unite  to  form  the  slightly  ionized,  molecular  hydroxide  that  sets 
the  other  equilibria  in  motion.  The  principle  at  the  basis  of  the 
change  is  thus  the  same  as  in  neutralization  (p.  398). 

In  ammonia,  nitrogen  is  trivalent,  while  in  the  salts  it  appears  to 
be  quinquivalent: 

H\  H\     /H  H\     /H 

H^N       H^N(  H-^N' 

n'         TL'      OH       w      ci 

HYDRAZINE,  HYDRAZOIC  ACID,  HYDROXYLAMINE 

Hydrazine  N2H^.  —  By  reduction  of  a  compound  of  nitric  oxide 
and  potassium  sulphite  by  means  of  sodium  amalgam,*  a  solution 
of  hydrazine  hydrate  is  obtained : 

K2S03,2NO  +  3H2  ->  N2H4,H20  +  K2S04. 

The  same  substance  is  more  easily  made  from  certain  organic  deriva- 
tives. When  the  hydrate  is  distilled  with  barium  oxide,  under 
reduced  pressure,  hydrazine  is  liberated: 

N2H4,H20  -f  BaO  -»  N2H4|  +  Ba(OH)2. 

Hydrazine  is  a  white  solid,  which  fumes  in  moist  air,  giving  the 
hydrate  once  more.  It  melts  at  1.4°  and  boils  at  113.5°. 

Hydrazine  hydrate  N2H5OH  freezes  at  about  —40°,  boils  at 
118.5°,  and  can  be  distilled  without  decomposition.  Its  aqueous 
solution  is  alkaline,  and  salts  are  formed  by  neutralization. 

Hydrazoic  Acid  HN%.  —  When  nitrous  oxide  (q.v.)  is  led  over 
sodamide  at  200°,  water  is  liberated  and  sodium  hydrazoate  remains 
behind : 

NH2Na  +  N2O  -*  NaN3  +  H20. 

A  dilute  solution  of  the  free  acid  is  best  obtained  by  distilling  the  lead 

*  The  sodium  dissolved  in  the  mercury  interacts  with  the  water,  giving  hydro' 
gen  (see  Active  state  of  hydrogen,  below). 


522  INORGANIC  CHEMISTRY 

salt  with  dilute  sulphuric  acid.  A  similar  solution  may  be  made  more 
directly  by  adding  cold  nitrous  acid  (q.v.)  to  a  cold  aqueous  solution 
of  hydrazine  hydrate: 

N2H5OH  +  HN02  -»  HN3  +  3H2O. 

By  repeated  distillation  of  the  solution  the  pure  acid  is  obtainable.  It 
boils  at  37°.  The  operation  is  a  dangerous  one,  as  the  pure  acid  is 
violently  explosive,  resolving  itself  into  nitrogen  and  hydrogen  with 
liberation  of  much  heat:  HN3,  Aq  -*  H  +  3N  +  Aq  +  61,600  cal. 
The  acid  has  an  odor  resembling  hydrochloric  acid,  and  the  sodium 
salt  has  a  taste  like  sodium  chloride.  The  silver  salt  is  insoluble  in 
water.  Thus  the  radical  N3  confers  properties  similar  to  the  radical 
Cl.  The  acid  has  a  somewhat  greater  activity  and  degree  of  ioniza- 
tion  than  has  acetic  acid.  Active  metals,  like  magnesium,  displace 
hydrogen  from  its  solution.  It  neutralizes  ammonium  hydroxide 
and  hydrazine  hydrate,  giving  two  salts,  NH4N3  and  N2H6N3.  These 
constitute  two  additional  compounds  of  nitrogen  and  hydrogen,  but 
differ  from  ammonia  and  hydrazine  in  being  ionogens.  Lead  hydra- 
zoate  Pb(N3)2  is  replacing  mercury  fulminate  in  the  manufacture  of 
percussion  caps. 

Hydroxylamine  NHZO.  —  Tin  displaces  hydrogen  from  dilute 
hydrochloric  acid :  Sn  -f  2HC1  — >  SnCl2  -f-  H2,  and  this  combination 
forms  a  reducing  agent  (see  below).  When  dilute  nitric  acid  (q.v.) 
is  added  to  the  mixture,  a  considerable  part  of  it  is  reduced  to  hydrox- 
ylamine: 

HN03  +  3H2  ->  NH3O  +  2H2O. 

The  hydroxylamine  forms  with  water  a  weak  base,  NH^O.OH,  which 
interacts  with  the  excess  of  acid,  giving  hydroxylamine  hydrochloride 
NH4OC1.  By  more  complete  reduction  of  part  of  the  nitric  acid, 
some  ammonium  chloride  NH4C1  is  formed  at  the  same  time.  To 
secure  the  salt  of  hydroxylamine,  the  tin  ions  are  removed  by  means 
of  hydrogen  sulphide,  which  precipitates  stannous  sulphide.  The 
filtered  solution  is  then  evaporated  to  dryness,  the  hydroxylamine 
hydrochloride  is  extracted  from  the  residue  with  absolute  alcohol, 
and  this  alcoholic  solution  is  finally  evaporated  in  turn.  The  hydro- 
chloride  is  a  white  crystalline  salt. 

A  better  yield  (80  per  cent)  is  obtained  by  placing  50  per  cent 
sulphuric  acid  (or  25  per  cent  HC1)  in  an  electrolytic  cell.  The 
negative  electrode  is  a  pool  of  mercury,  or  a  small  piece  of  lead 


NITROGEN  AND  ITS  COMPOUNDS  WITH  HYDROGEN      523 

amalgamated  with  mercury.  Hydrogen  is  liberated  on  the  surface 
of  the  mercury,  which  acts  as  a  contact  agent,  an'd  nitric  acid  is 
allowed  to  flow  very  slowly  from  a  tube  on  to  the  surface  of  the 
mercury.  The  cell  is  kept  cool,  and  a  high  current  density,  by  using 
a  small  mercury  surface,  is  required  for  a  good  yield.  The  equation 
is  the  same  as  that  given  above. 

Anhydrous  hydroxylamine  is  best  obtained  by  heating  the  ortho- 
phosphate  (NH4O)3P04  under  reduced  pressure: 

(NH40)3P04  -»  H3P04  +  3NH30  T  - 

It  is  a  white  solid  melting  at  33°  and  boiling  at  58°  at  22  mm.  pressure. 
Even  before  melting  (above  15°),  it  begins  to  decompose,  and  explodes 
at  or  below  130°.  In  chemical  behavior  it  is  like  ammonia.  With 
water  it  forms  a  base  which  combines  with  acids,  but  is  less  active 
than  ammonium  hydroxide.  It  is  a  stronger  reducing  agent  than 
ammonia,  precipitating  silver  from  a  solution  of  silver  nitrate. 

The  union  with  acids  indicates  that  the  molecule  of  hydroxyl- 
amine is  unsaturated,  and  hence  the  nitrogen  unit  is  supposed  to  be 
trivalent : 

H\  H\         H 

H-N  H-N( 

H-0/  H-0/         Cl 

Halogen  Compounds  of  Nitrogen.  — -  When  ammonium  chlo- 
ride solution  is  treated  with  excess  of  chlorine,  drops  of  an  oily 
liquid,  nitrogen  trichloride  NC13,  are  formed : 

3C12  +  3H2O  <=*  3HC1  +  3HOC1  +  NH4C1  ->  NC13  +  3H20  +  4HC1. 

It  is  extremely  explosive,  resolving  itself  into  its  constituents  with 
.liberation  of  much  heat. 

When  a  solution  of  iodine  in  potassium  iodide  solution  (p.  276) 
is  added  to  aqueous  ammonia,  a  brown  precipitate  is  formed.  The 
composition  of  the  substance  depends  upon  the  temperature.  Thus, 
at  -60°  it  is  NI3,12NH3,  at  -40°  NI^NHg,  at  -35°  NI3,2NH3,  and 
at  —25°  and  above  NI3,NH3.  The  first  three,  when  the  temperature 
rises,  lose  ammonia,  but  the  last,  commonly  named  nitrogen  iodide, 
explodes.  It  may  be  handled  while  wet,  but  when  dry  decomposes 
into  its  constituents  with  violent  explosion  if  touched  with  a  feather. 
The  ammonia  it  contains  is  like  that  in  CaCl2,8NH3  (p.  517),  and 
resembles  water  of  hydration. 


524  INORGANIC  CHEMISTRY 

Exercises.  —  1.  When  moist  air  is  used  as  a  source  of  nitrogen, 
what  advantage's  there  in  using  copper  rather  than  the  less  expensive 
metal  iron,  for  removing  the  oxygen  (p.  514)? 

2.  How  many  grams  of  water  at  0°  could  be  frozen  (p.  144)  by  the 
removal  of  the  heat  required  to  evaporate  50  g.  of  liquid  ammonia 
(p.  517)? 

3.  How  many  grams  of  ammonia  are  contained  in  1 1.  of  "concen- 
trated ammonia"  (p.  517)?     What  would  be  the  change  in  volume, 
if  this  were  made  by  mixing  liquid  ammonia  with  water? 

4.  What  are  the  ions  of  hydrazine  hydrate  (p.  521)?     Formulate 
(p.  387)  the  neutralization  of  this  base  with  sulphuric  acid. 

5.  What  is  the  object  attained  by  distilling  under  reduced  pressure 
in  making  hydrazine  (p.  521)  and  hydroxylamine  (p.  522)? 

6.  Classify  (p.  2?8)  the  interaction  of  a  nitride  with  water  (p.  514), 
and  of  chlorine- water  and  ammonium  chloride  (p.  523),  and  the 
results  of  heating  ammonium  nitrite  (p.  514)  and  ammonium  chloride 
(p.  520). 

7.  What  substances  are  present  in  ammonium  hydroxide  solution? 
When  the  liquid  is  heated,  what  happens  to  each?     Formulate  the 
system. 


CHAPTER  XXVI 
OXIDES  AND  OXYGEN  ACIDS   OP  NITROGEN 

THE  names  and  formulae  of  the  oxides  and  oxygen  acids  of  nitrogen 
are  as  follows: 

Nitrous  oxide  N2O          < Hyponitrous  acid  H2N202 

Nitric  oxide  NO 

Nitrous  anhydride  N2O3  < >      Nitrous  acid  HN(>2 

Nitrogen  tetroxide  N204  and  N02 

Nitric  anhydride  N205     < >      Nitric  acid  HN03. 

All  the  oxides  are  endothermal  compounds,  yet,  with  the  exceptions 
of  the  third  and  the  last,  they  are  all  relatively  stable.  The  acids, 
when  deprived  of  the  elements  of  water,  yield  the  oxides  opposite 
which  they  stand.  Conversely,  excepting  in  the  case  of  nitrous  oxide, 
the  anhydrides  with  water  give  the  acids.  All  of  these  substances  are 
obtained  directly  or  indirectly  from  nitric  acid  —  nitric  anhydride  by 
removal  of  water,  the  others  by  reduction.  We  turn,  therefore,  first, 
to  nitric  acid,  its  sources  and  properties. 

NITRIC  ACID 

Sources.  —  Sodium  nitrate,  or  Chile  saltpeter  (caliche),  is  found 
in  a  desert  region  near  the  boundary  of  Chile  and  Peru,  and  chiefly  in 
the  former  country.  The  deposit  is  about  5  feet  thick,  2  miles  wide, 
220  miles  in  length,  and  contains  20  to  60  per  cent  of  the  salt.  Puri- 
fication is  effected  by  recrystallization.  Potassium  nitrate,  or  Bengal 
saltpeter,  is  found  in  the  soil  in  the  neighborhood  of  cities  in  India, 
Persia,  and  other  oriental  countries.  It  arises  from  the  oxidation  of 
animal  refuse  (cf.  p.  517)  through  the  mediation  of  nitrifying  bacteria. 
The  potash  and  lime  in  the  soil,  along  with  the  product  of  oxidation  of 
the  nitrogen,  give  nitrates  of  potassium  and  calcium.  The  aqueous 
extract  of  this  soil  is  treated  with  wood  ashes,  which  contain  potash 
K2C03.  It  is  poured  off  from  the  calcium  carbonate  thus  precipitated, 
and  is  finally  evaporated.  The  organic  compounds  of  nitrogen 
originally  contained  in  guano,  a  valuable  fertilizer,  are  frequently 
found  to  have  been  changed  into  nitrates  by  nitrifying  bacteria. 

525 


526  INORGANIC  CHEMISTRY 

The  action  of  the  nitrifying  bacteria  may  be  imitated  in  a  rough  way 
[Lect.  exp.].  Air  is  caused  to  pass  slowly  through  concentrated  aqueous  am- 
monia, whereby  it  becomes  mixed  with  ammonia  gas.  This  mixture  is  led 
through  a  wide  tube  containing  platinized  asbestos  and  is  then  discharged  into 
a  large  flask.  When  the  asbestos  is  warmed,  it  begins  to  glow,  and  thereafter 
the  action  maintains  itself.  A  part  of  the  ammonia  is  oxidized  to  nitric  acid, 
which  combines  with  the  excess  of  ammonia,  giving  ammonium  nitrate.  This 
salt  forms  a  cloud  of  solid  particles  which  settle  in  the  flask.  This  process  is 
used  on  a  commercial  scale. 

Manufacture. — When  any  nitrate  is  treated  with  any  acid,  nitric 
acid  is  formed  by  a  reversible  double  decomposition.  As  sodium 
nitrate  is  the  cheapest  salt  of  nitric  acid,  it  is  always  employed.  For 
the  same  reason,  and  on  account  of  its  activity  and,  above  all,  because 
of  its  relative  involatility,  sulphuric  acid  is  used  to  displace  it: 

NaN03  +  H2S04  *=?  NaHS04  +  HNO3  T- 

The  nitric  acid  is  rather  volatile  (b.-p.  86°),  while  sulphuric  acid  (b.-p. 
330°)  is  much  less  so,  and  the  two  salts  are  not  volatile  at  all.  Thus 
the  interaction  proceeds  to  completion  very  easily  (cf.  p.  207,  see  also 
p.  559).  The  materials  are  heated  in  cast-iron  stills,  and  the  vapor 
is  condensed  in  glass  or  earthenware  pipes  surrounded  by  water.  In 
many  factories  a  reduced  pressure  is  maintained  in  the  stills  and 
condensers,  in  order  that  the  distillation  may  take  place  at  the  lowest 
possible  temperature.  This  precaution  is  taken  to  reduce  to  a 
minimum  the  partial  decomposition  of  the  nitric  acid  (see  below). 

Physical  Properties.  —  Nitric  acid  is  a  colorless,  mobile  liquid 
(density  1.52)  boiling  at  86°,  and  freezing  to  a  solid  (m.-p.  —47°). 
It  fumes  strongly  when  its  vapor  issues  into  moist  air  (cf.  p.  211). 
An  aqueous  solution  containing  68  per  cent  of  the  acid  boils  at  120.5°, 
while  the  pure  acid,  pure  water,  and  all  other  mixtures,  boil  at  lower 
temperatures,  and  have,  therefore,  higher  vapor  pressures.  On  this 
account  a  more  dilute  acid,  when  heated,  loses  water  until  it  reaches 
this  strength.  The  68  per  cent  nitric  acid  of  constant  boiling-point 
(p.  211)  forms  the  " concentrated  nitric  acid"  of  commerce  (density 
1.41).  • 

Chemical  Properties.  —  1.  Like  chloric  acid  (p.  482),  and  other 
oxygen  acids  of  the  halogens,  nitric  acid  is  most  stable  when  mixed 
with  water.  The  pure  (100  per  cent)  acid  decomposes  while  being 
distilled: 

4HN03  -» 4N02  +  2H20  +  02, 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  527 

yet  not  with  explosive  violence  like  chloric  acid.  The  distillate  is 
colored  brown  by  dissolved  nitrogen  tetroxide  NO2.  Repeated  dis- 
tillation finally  leaves  68  per  cent  of  the  acid,  mixed  with  32  per  cent 
of  water  formed  by  the  above  decomposition.  The  acid  of  constant 
boiling-point  is,  therefore,  reached,  as  usual,  from  more  concentrated 
as  well  as  from  less  concentrated  specimens. 

" Fuming"  nitric  acid  is  brown  in  color,  and  contains  a  con- 
siderable amount  of  dissolved  nitrogen  tetroxide.  It  is  made  by 
distilling  the  acid  with  a  little  starch.  The  latter  reduces  a  part  of 
the  nitric  acid  and  liberates  more  of  the  tetroxide  than  does  mere 
distillation. 

2.  Nitric  acid,  when  dissolved  in  water,  is  highly  ionized,  and 
is  therefore  active  as  an  acid.     By  interaction  with  hydroxides  and 
oxides  it  forms  nitrates. 

3.  When  pure  nitric  acid  (b.-p.  86°)  is  poured  upon  phosphoric 
anhydride,  the  latter  combines  with  the  elements  of  water,  and  dis- 
tillation gives  nitric  anhydride:  2HNO3  +  P205  ->  N2O5  T  +  2HP03. 
The  anhydride  is  a  white  solid  melting  at  30°  and  boiling  at  45°. 
It  unites  vigorously  with  water  to  form  nitric  acid.     It  cannot  be 
kept,  as  it  decomposes  into  nitrogen  tetroxide  and  oxygen,  2N2Os  — > 
4NO2  +  02,  with  liberation  of  heat. 

4.  Like  the  unstable  oxygen  acids  of  the  halogens,  nitric  acid  is  an 
oxidizing  agent  even  when  diluted  with  water.     The  multiplicity  of 
the  products  into  which  it  may  be  decomposed  by  reduction,  however, 
renders  separate  treatment  of  this  property  necessary  (see  p.  534). 

5.  Nitric    acid   interacts  energetically  with   many  compounds 
of  carbon  to  give  nitro-derivatives.     Thus,  when  heated  with  phe- 
nol C6H5OH    (carbolic  acid)   it  gives   picric   acid   (trinitrophenol) 
C6H2(NO2)3OH,  which  crystallizes  in  yellow  needles  in  the  mixture. 
This  is  a  yellow  dye,  used  also  as  an  explosive  (lyddite) : 

C6H6(OH)  +  3HON02  ->  C6H2(OH)(NO2)3  +  3H20. 

The  presence  of  water  decreases  the  activity  of  the  molecules.  Hence, 
in  this  sort  of  action,  which  is  not  ionic,  not  only  is  concentrated 
nitric  acid  employed,  but  concentrated  sulphuric  acid  is  added  to 
assist  in  the  elimination  of  the  water  (cf.  p.  438).  Nitric  acid,  heated 
with  toluene  CeHsCHs,  gives  trinitrotoluene: 

CH3C6H5  +  3HON02  ->  CH3C6H2(N02)3  +  3H20. 

This  substance  (T.N.T.)  is  used  for  filling  "high  explosive"  shells, 
because  it  can  be  melted  (m,-p,  81,5°)  and  poured  in,  making  the 


528  INORGANIC   CHEMISTRY 

filling  easy,  safe,  rapid,  and  complete.  It  is  not  easily  exploded  by 
shocks  during  transportation,  but  it  explodes  instantaneously  and 
completely  with  a  detonator.  The  following  equation  shows, 
roughly,  the  decomposition,  and  the  large  amount  of  carbon  set 
free  explains  the  black  smoke  produced: 

2CH3C6H2(N02)3  ->  5H20  +  3N2  +  7CO  +  70. 

It  will  be  seen  that  the  group  NO2  has  taken  the  place  of  hydro- 
gen, which  was  formerly  attached  directly  to  the  carbon  of  the  phenol 
or  toluene.  Compounds  of  this  kind  are  called  nitro-derivatives. 

6.  Organic    compounds    of   another    class,    the   alcohols    (q.v.), 
interact  with  molecular  nitric  acid  in  a  different  way.     The  latter  is 
mixed  with  sulphuric  acid  with  the  same  object  as  before.     Thus, 
when  glycerin  is  added  slowly  to  the  cooled  mixture,  glyceryl  nitrate 
(so-called  nitro-glyceiine,  see  below)  is  produced : 

C3H6(OH)3  +  3HON02  -4*  C3H6(ON02)3  +  3H20. 

Here  it  is  the  hydrogen  of  the  hydroxyl  groups  that  is  displaced  by 
N02.  The  action  is  not  ionic,  and  the  product  is  not  an  ionogen. 
Gun-cotton  is  made  by  this  action,  cotton  (cellulose)  being  employed : 

(C6H1005)2  +  6HON02  -?  C12H1404(ON02)6  +  6H20. 

7.  Nitric  acid  produces  substances  of  bright-yellow  color,  known 
as  xanthoproteic  acids,  when  it  comes  in  contact  with  proteins,  e.g., 
in  the  skin,  or  in  wool.     Hence  nitric  acid  stains  woolen  clothing 
yellow.     This  reaction  is  used  as  a  test  for  proteins. 

The  chemical  properties  of  nitric  acid  are  best  represented  by  the 
graphic  formula  (see  p.  540) : 

*0 
H-  0  -  N 


8.  Nitron  (1,4-diphenyl-endanilino-dihydrotriazole)  C2oHi6N4  gives 
a  fairly  insoluble  nitrate  C2oHieN4,HN03,  when  nitron  acetate  is  added 
to  a  solution  containing  nitric  acid.  Nitric  acid  can  be  determined 
quantitatively  by  weighing  the  precipitate. 

Nitrates.  —  The  nitrates  of  the  metallic  elements  are  all  more  or 
less  easily  soluble  in  water.  When  heated  they  decompose  in  one  or 
other  of  three  ways  (see  pp.  531,  537,  539).  Sodium  nitrate  is 
used  largely  as  a  fertilizer.  Much  is  employed  in  sulphuric  acid 
manufacture,  and  the  rest  for  conversion  into  potassium  nitrate  and 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  529 

in  making  nitric  acid.  Potassium  nitrate  is  used,  along  with  sulphur 
and  charcoal,  in  the  manufacture  of  gunpowder.  The  individual 
nitrates  are  described  under  the  metallic  elements  they  contain. 

NITRIC  OXIDE  AND  NITROGEN  TETROXIDE 

Preparation  of  Nitric  Oxide  NO.  —  Pure  nitric  oxide  is  ob- 
tained by  adding  nitric  acid  to  a  boiling  solution  of  ferrous  sulphate 
in  dilute  sulphuric  acid,  or  of  ferrous  chloride  in  hydrochloric  acid: 

2FeS04  +  H2S04  ->  Fe2(SO4)3  (+  2H)    X  3          (1) 

(3H)  +  HNO3  ->  NO  +  2H2O  X  2  (2) 

6FeSO4  +  3H2SO4  +  2HNO3  -»  3Fe2(SO4)3  +  2NO  +  4H2O 

The  first  partial  equation  does  not  take  place  at  all  unless  an  oxidizing 
agent  like  nitric  acid  is  present  (p.  320).  The  multiplication  of  the 
two  partial  equations  by  3  and  2,  respectively,  is  required  in  order 
that  the  hydrogen,  which  is  not  a  product,  may  cancel  out.  This 
action  is  used  as  a  means  of  determining  the  quantity  of  nitric  acid  in 
a  solution,  or  of  nitrates  in  a  mixture,  by  measurement  of  the  volume 
of  nitric  oxide  evolved. 

Nitric  oxide  may  also  be  obtained  when  sufficiently  dilute  nitric 
acid  (density  1.2)  acts  upon  copper  (see  p.  535).  Although  some 
nitrous  oxide  and  nitrogen  are  formed,  this  interaction  furnishes 
a  more  convenient  method  of  generating  the  gas  in  the  laboratory 
(see  also  p.  532). 

Properties  of  Nitric  Oxide.  —  Nitric  oxide  is  a  colorless  gas. 
In  solid  form  it  melts  at  —167°  and  the  liquid  boils  at  —153.6°. 
Its  solubility  in  water  is  slight.  The  density  of  the  gas  shows  the 
formula  to  be  NO ;  and  there  is  no  tendency  to  form  a  polymer,  such 
as  N2O2,  even  at  low  temperatures. 

Nitric  oxide  is  the  most  stable  of  the  oxides  of  nitrogen.  Vig- 
orously burning  phosphorus  continues  to  burn  in  the  gas,  the  heat 
evolved  liberating  the  oxygen  required  for  the  continuation  of  the 
combustion.  Burning  sulphur  and  an  ignited  taper,  however, 
cannot  decompose  the  gas,  and  are  extinguished. 

Nitric  oxide  has  two  characteristic  properties.  It  unites  directly 
with  oxygen  in  the  cold  to  form  the  reddish-brown  nitrogen  tetroxide : 

2NO  +  O2  <=»  2N02. 
This  can  be  used  as  a  test  for  even  a  trace  of  free  oxygen,  mixed  with 


530  INORGANIC  CHEMISTRY 

other  gases.     The  same  result  follows  when  it  is  led  into  warm  con- 
centrated nitric  acid:  NO  +  2HNO3  ?±  3NO2  +  H20. 

It  also  unites  with  a  number  of  salts,  the  compound 
in  the  case  of  ferrous  sulphate,  FeNO.S04  (a  molecular 
compound,  see  below),  being  capable  of  existence  in 
solution  and  possessing  a  brown  color.  The  NO  is 
part  of  the  positive  ion  FeNO"1"1",  and  migrates  with 
it  (Manchot). 

Since  ferrous  sulphate  will  first  reduce  nitric  acid  to 
nitric  oxide  (p.  529),  and  the  excess  of  the  salt  will  then 
give  a  brown  color  with  the  product,  a  delicate  test  for 
nitric  acid  is  founded  upon  the  above  action.     The  sub- 
stance supposed  to  contain  a  nitrate  is  mixed  with  a 
strong  solution  of  ferrous  sulphate,  and  concentrated 
sulphuric  acid  is  poured  down  the  side  of  the  tube  so  as 
to  lie  below  the  lighter  mixture   (Fig.   122).     At  the 
surface   of   contact   the   sulphuric    acid   liberates   the 
FIQ  122        nitric  acid,  this  is  reduced  to  nitric  oxide,  the  molecular 
compound  is  formed,  and  a  brown  layer  is  seen.     Even 
when  the  amount  of  the  nitrate  is  very  small,  the  brown  tint  is 
distinctly  visible,  by  contrast  with  the  colorless  liquids  above  and 
below  it. 

Molecular  Compounds.  —  When  substances  formed  by  union 
of  two  compounds  have  a  prevailing  tendency  to  decompose  into  the 
same  two  materials,  and  exhibit  the  chemical  properties  of  their  con- 
stituents rather  than  individual  ones  of  their  own,  they  are  often 
called  molecular  compounds.  Thus  the  above  substance  FeNO.SO4 
gives  off  the  nitric  oxide  again  when  warmed.  Similarly,  hydrates 
(pp.  150-154)  are  formed  by  union  of  salts  or  other  substances  with 
water,  and  are,  for  the  most  part,  decomposed  by  solution.  Double 
salts  (p.  402),  such  as  ferrous-ammonium  sulphate  FeSO4,(NH4)2- 
S04,6H2O,  of  which  very  many  are  known,  are  of  the  same  character. 
They  are  stable  only  in  the  solid  form.  There  are  also  compounds  of 
salts  with  ammonia  (see  Copper  and  silver),  and  with  carbon  monoxide 
CO,  one  such  compound  being  formed  with  cuprous  chloride  (q.v.). 

The  name  molecular  compounds  is  derived  from  the  supposition 
that,  in  these  compounds,  the  molecules  of  the  components  retain 
their  integrity  to  some  extent  and  are  thus  ready  to  be  liberated. 
This  is  an  attempt  to  explain  the  fact  that  the  behavior  is  that  of  the 
constituents.  It  distinguishes  molecular  compounds  from  substances 


OXIDES  AND  OXYGEN   ACIDS  OF  NITROGEN  531 

like  ammonium  chloride  and  phosphorus  pentachloride.  The  former 
may  be  made  by  union  of  HC1  and  NH3,  but  usually  behaves  rather 
as  if  composed  of  NH4  and  Cl.  The  latter  (q.v.)  dissociates  into  PC13 
and  Cl2,  but  with  water  gives  phosphoric  acid  (cf.  p.  210),  which  is 
derivable  from  the  pentachloride  only.  The  distinction  is  of  practical 
rather  than  theoretical  importance,  however,  for  there  are  all  grada- 
tions in  the  behavior  of  molecular  compounds.  It  is  useful  simply 
as  a  rough  means  of  classifying  and  remembering  certain  facts. 

Distinguishing  molecular  compounds  from  ordinary  compounds  is 
further  justified  by  the  fact  that  the  constituents  of  molecular  com- 
pounds often  seem  to  be  saturated  (p.  427) ,  and  no  ordinary  valences 
are  available  for  holding  the  new  material.  Thus  in  Ca^Cla1  the 
ordinary  valences  are  all  saturated.  Yet  the  salt  forms  the  hydrate 
CaCl2,6H2O  with  water  H2I011,  which  is  likewise  a  saturated  compound. 
The  conception  of  molecular  compounds  implies,  therefore,  the  idea 
of  a  sort  of  valence  of  molecules.  Thus  FeS04  forms  FeS04,7H2O 
and  FeSO4,(NH4)2SO4,6H2O,  and  FeSO4,K2SO4,6H2O,  in  all  of  which 
seven  other  molecules  are  combined  with  it.  The  sulphates  of 
other  bivalent  metals  (q.v.),  such  as  copper  and  magnesium,  form 
molecular  compounds  of  the  same  nature.  Ammonium  chloride,  on 
the  other  hand,  is  not  a  molecular  compound,  because,  although  NH3 
unites  with  HC1,  HBr,  HI,  and  HF,  yet  nitrogen  is  quinquivalent, 
and  substances  like  N2O5,  NH4C1,  etc.,  may  fitly  be  regarded  as 
ordinary  compounds. 

Preparation  of  Nitrogen  Tetroxide  N2O4  and  NOZ.  —  This 
substance  is  liberated  by  heating  nitrates,  other  than  those  of  po- 
tassium, sodium,  or  ammonium,  such  as  the  nitrates  of  lead  and 
copper: 

2Cu(NO3)2  -rt  2CuO  +  4N02  +  O2. 

The  nitrates  of  metals  above  mercury  in  the  E.  M.  series  (p.  404) 
leave  the  oxide.  The  nitrates  of  mercury  and  of  the  metals  below  it 
leave  the  metal.  When  the  mixed  gases  are  led  through  a  U-tube 
immersed  in  a  freezing  mixture,  the  tetroxide  condenses  as  a  pale- 
yellow  liquid  (b.-p.  22°,  m.-p.  —10.5°),  and  the  oxygen  passes  on. 

The  compound  may  also  be  made  by  direct  union  of  nitric  oxide 
and  oxygen,  or  by  oxidation  of  nitric  oxide  by  concentrated  nitric 
acid  (p.  530).  It  is  likewise  almost  the  sole  product  of  the  inter- 
action of  concentrated  nitric  acid  and  copper  (see  p.  535).  If  any 
nitric  oxide  were  produced  by  the  primary  action,  it  would  be  oxidized 
to  nitrogen  tetroxide  in  passing  up  through  the  concentrated  acid. 


532  INORGANIC  CHEMISTRY 

Properties  of  Nitrogen  Tetroxide.  —  The  most  striking  pecu- 
liarity of  this  gas  is  that,  when  hot,  it  is  deep  brown  in  color,  and 
when  cold,  pale  yellow.  The  density  of  the  gas  decreases  very 
rapidly  from  22°  to  140°,  and  increases  again  as  the  temperature  falls. 
The  molecular  weights  calculated  from  these  observations  are:  'at 
22°,  90;  at  70°,  55.6;  at  135°,  46.3;  at  154°,  45.7.  Now  the  molecular 
weights  corresponding  to  the  formulae  N204  and  NO2  are  92  and  46, 
respectively,  so  that  these  results  mean  that  the  deep-brown  gas  is 
N02,  and  that  as  this  is  cooled  it  combines  to  form  the  colorless  N2O4. 
Measurement  of  the  depression  the  substance  causes  in  the  freezing- 
point  (cf.  p.  336)  of  glacial  acetic  acid  gives  the  molecular  weight  92, 
so  that  in  solution  and  at  the  temperature  of  freezing  acetic  acid 
(below  17°)  the  substance  is  all  N2O4. 

When  the  temperature  is  carried  above  154°,  by  passing  the 
brown  gas  through  a  red-hot  tube,  the  brown  color  disappears  once 
more,  and  nitric  oxide  and  oxygen  are  formed.  On  cooling,  the  same 
steps  through  brown  gas  to  pale-yellow  gas  are  retraced  : 


Colorless  Brown         Colorless 

Since  nitrogen  tetroxide  yields  free  oxygen  more  readily  than  does 
nitric  oxide,  phosphorus  burns  readily  in  it;  a  taper,  or  burning 
sulphur,  however,  is  extinguished.  It  has  powerful  oxidizing  proper- 
ties, and  "fuming  nitric  acid,"  which  contains  it  in  solution,  is 
employed  when  oxidation  is  the  special  object  in  view.  For  the 
same  reason,  the  gas  is  sometimes  used  in  bleaching  flour. 

This  oxide  is  intermediate  in  composition  between  nitrous  and 
nitric  anhydrides,  and,  when  dissolved  in  cold  water,  gives  both  nitric 
and  nitrous  acids:  N2O4  +  H20  —  >  HNO3  +  HN02.  If  a  base  is 
present,  a  mixture  of  the  nitrate  and  nitrite  of  the  metal  is  produced 
(cf.  p.  474).  When  the  water  is  not  cooled,  the  nitrous  acid  (q.v.), 
being  unstable,  gives  nitric  oxide  and  nitric  acid,  so  that  the  result  is  : 
3N02  +  H20  <=>  2HN03  +  NO. 

NITEIC  ACID  FROM  ATMOSPHERIC  NITROGEN 

The  Reactions  Involved.  —  Nitrogen  and  oxygen  have  no 
tendency  to  unite  at  room  temperature  to  form  nitric  oxide.  The 
union  is  endothermal,  and  is  therefore  favored  by  a  high  temper- 
ature (van't  HofFs  law,  p.  305)  : 

N2  +  02  +  43,200  cal.  <=±  2NO. 
Even  at  1922°,  however,  using  atmospheric  air,  only  1  per  cent 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN 


533 


nitric  oxide  is  formed,  and  at  2927°,  5  per  cent.  The  electric  dis- 
charge actually  used  gives  about  1  per  cent. 

The  mixture  is  next  cooled,  to  permit  the  union  of  2NO  +  02 
<=^2N02,  because  (p.  532)  nitrogen  tetroxide  is  decomposed  at 
about  154°,  and  therefore  cannot  be  formed  at  1900°. 

Next,  the  air  containing  NO2  is  passed  through  absorbing  towers, 
down  which  water  trickles,  and  nitric  acid  is  formed : 

3N02  +  H20  ->  2HN03  +  NO. 

The   NO  liberated  combines  with  more  atmos-      *c™vT'| 
pheric    oxygen    to    form    NO2,    which    interacts 
again  with  the  water,  and  practically  no  nitric 
oxide  is  lost. 

Finally,  the  nitric  acid  is  poured  upon  lime- 
stone CaC03,  and  the  calcium  nitrate  formed  is 
sold,  for  use  as  a  fertilizer,  under  the  name  air 
saltpeter. 

The    Plant    used   in    the    Fixation.  —  At 

Notodden  and  elsewhere  in  Norway,  the  Birke- 
land-Eyde  process  (Fig.  123)  is  used.  Hydro- 
electric power  is 
employed,  and  an 
arc  discharge  be- 
tween two  rods  of 
carbon  is  spread, 
by  the  influence  of 
large  and  powerful 
electromagnets,  in- 
to a  circular  brush 
discharge  several 
feet  in  diameter. 
The  figure  is  a  cross 
section  of  the  space 
filled  by  the  dis- 
FlG  123.  charge,  the  small 

circle  in  the  center 
being  a  section  of  one  carbon  rod.  Air  is  blown  through  the  flame 
in  such  a  way  that  none  can  avoid  passing  through  at  least  a  part  of 
the  heated  area.  The  yield  is  about  70  g.  of  nitric  acid  per  kilowatt- 
hour,  and  the  net  earnings  are  $350,000  (1911). 


Fia.  124. 


534  INORGANIC  CHEMISTRY 

The  Badische  process,  used  in  the  same  factories  in  Norway, 
employs  a  discharge  through  a  tube  over  20  feet  long  (Fig.  124). 
The  stream  of  air  rotates  as  it  traverses  the  tube,  so  that  every  part 
is  exposed  to  the  discharge.  The  Pauling  process,  used  at  Gelsen- 
kirchen  in  Germany  and  Nitrolee,  South  Carolina,  uses  preheated 
air  and  a  different  arrangement  of  the  discharge. 

Other  reactions  involving  the  fixation  of  atmospheric  nitrogen 
are  discussed  under  cyanamide  (q.v.)  and  root  nodules  (p.  515). 

OXIDIZING  ACTIONS  OF  NITRIC  ACID 

When  nitric  acid  gives  up  oxygen  to  any  body,  it  is  itself  reduced. 
Hence,  according  to  convenience,  we  shall  refer  to  oxidations  by,  or 
reductions  of  nitric  acid. 

Oxidation  of  Hydrogen.  —  The  metals  preceding  hydrogen  in 
the  electromotive  series  (p.  404)  displace  hydrogen  from  nitric 
acid,  as  they  do  from  other  acids.  With  metals  more  active  than 
zinc,  such  as  magnesium,  a  great  part  of  the  hydrogen  escapes  in 
the  free  condition.  But,  in  the  case  of  zinc  and  the  metals  below  it, 
most  or  all  of  the  hydrogen  is  oxidized  to  water  by  the  nitric  acid, 
and  part  of  the  acid  is  reduced  (see  Active  hydrogen,  p.  543) .  Thus, 
with  zinc  and  very  dilute  nitric  acid,  almost  the  only  product,  aside 
from  zinc  nitrate,  is  ammonia: 

4Zn+    8HN03  -*  4Zn(NO3)2  (+  8H)  (1) 

(8H)  H-      HNO3  -»  NH3  +  3H20  (2) 

NH3  +      HNO3  ->  NH4NO3 (3) 

4Zn  +  10HNO3  -5  4Zn(N03)2  +  NH4NO3  +  3H2O 

With  the  excess  of  nitric  acid,  ammonium  nitrate  is  formed. 

Heavy  Metals.  —  The  less  active  metals,  such  as  copper  and 
silver,  do  not  displace  hydrogen  from  dilute  acids  (p.  129),  but 
reduce  nitric  acid,  nevertheless,  and  are  converted  into  nitrates. 
Platinum  and  gold  (cf.  p.  439)  alone  are  not  attacked.  Thus,  copper, 
with  somewhat  diluted  nitric  acid  (density  1.2)  gives  cupric  nitrate 
and  nitric  oxide  NO. 

In  making  the  equation  for  this  action  we  may  use  the  anhydride 
plan  (p.  496),  which  is  applicable  whenever  an  oxygen  acid  gives 
an  oxide  by  reduction.  We  resolve  the  formula  of  nitric  acid  into 
-nose  of  water  and  the  anhydride  H2O,N206  (=  2HNO3).  This 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  535 

shows  that  the  two  molecules  of  the  acid  will  give  2NO,  and  30  will 
remain : 

2HN03  ( =  H20,N205)  -»  H20  +  2NO  (+30)  (1) 

(3O)  +  6HN03  +  3Cu  -»  3H20  +  3Cu(N03)2 (2) 

8HNO3  +  3Cu  ->  4H20  +  2NO  +  3Cu(N03)2 

The  nitric  oxide  is  liberated  as  a  colorless  gas,  but  forms  the  brown 
tetroxide  at  once  on  meeting  the  oxygen  of  the  air  (p.  529). 

When  concentrated  nitric  acid  is  used  with  copper,  almost  pure 
nitrogen  tetroxide  is  obtained: 

2HN03  ( =  H20,N205)  -»  H20  +  2N02  (+  0)  (1) 

(O)  +  2HN03  +  Cu  ->  H20  +  Cu(N03)2  (2) 

4HNO3  +  Cu  -*  2H20  -f  2N02  +  Cu(NO3)2 

The  equations  for  actions  like  the  above  may  be  built  up  from  partial  equations 
of  various  kinds  (cf.  p.  269).  Thus  we  may  begin  by  forming  the  nitrate  of  the 
metal,  and  then  use  the  balance,  consisting  of  hydrogen,  along  with  other  mole- 
cules of  nitric  acid  to  secure  the  oxide  and  water: 

Cu  +  2HN03  -»  Cu(N03)2  (+  2H) 
(2H)  +  2HNO3  ->  2H2O  +  2NO2 


Cu  +  4HNO3  -» Cu(NO3)2  +  2H2O  +  2NO2 

As  the  subdivision  is  purely  arithmetical  (p.  270),  this  procedure  does  not  involve 
the  assumption  that  copper  does  actually  displace  hydrogen  as  a  free  element. 
Yet  it  would  not  necessarily  be  incorrect  to  make  even  this  supposition.  Although 
unable  to  liberate  hydrogen  in  quantity  from  a  dilute  acid,  copper  must  be  held 
to  displace  a  minute  amount  of  it: 

Cu  +  2HNO3  <=5  Cu(NO3)2  +  H2,      or      Cu  +  2H+  ±3  CU++  +  H2, 

and  to  be  restrained  by  the  much  more  vigorous  reverse  action  (p.  404)  from  con- 
tinuing this  operation.  In  this  point  of  view  the  oxidation  of  the  trace  of  free 
hydrogen  by  the  excess  of  nitric  acid  continuously  annihilates  the  possibility  of 
reverse  action. 

Complexities  of  Oxidation  by  Nitric  Acid.  —  The  above  are 
types  of  the  interactions  of  metals  with  nitric  acid.  In  actual  experi- 
ments the  behavior  is  usually  more  complex.  Thus,  as  a  rule,  the 
action  is  very  slow  at  first,  and  gathers  speed  with  the  accumulation 
of  the  reduction  products,  which  act  eatalytically. 

Again,  different  concentrations  of  nitric  acid  give  different  prod- 
ucts with  the  same  metal.  The  reader  should  note  the  constant 
production  of  nitric  oxide  with  diluted  acid,  and  the  invariable 


536  INORGANIC  CHEMISTRY 

formation  of  nitrogen  tetroxide  with  concentrated  acid.  This  is  ex- 
plained by  the  fact  that  nitrogen  tetroxide  cannot  pass  unchanged 
through  a  liquid  containing  much  water,  for  it  gives  nitric  acid  and 
nitric  oxide  with  the  latter  (p.  532).  Conversely,  where  the  nitric 
acid  is  concentrated,  nitric  oxide,  even  if  formed  by  the  interaction 
with  the  metal,  must  be  oxidized  to  nitrogen  tetroxide  as  it  passes  up 
through  the  liquid  (p.  530).  Note,  also,  that  the  nitrate  of  the 
metal  is  formed,  if  the  nitrate  is  stable,  not  the  oxide. 

Finally,  intermediate  concentrations  give  mixtures  of  these  two 
oxides,  and,  with  zinc,  even  nitrous  oxide  N2O  and  nitrogen  may  be 
found  in  considerable  quantities  in  the  gases  evolved. 

Oxidation  of  Non-Metals.  —  With  non-metals  the  actions  are 
different,  in  so  far  that  these  elements  form  no  nitrates.  Thus, 
sulphur  boiled  in  nitric  acid  gives  sulphuric  acid,  along  with  nitric 
oxide,  equation  (3),  or  with  nitrogen  tetroxide,  equation  (6),  or  with 
both,  according  to  the  concentration  of  the  acid  (see  above) : 

2HN03  ( =  H20,N205)  ->  2NO  +  H20  (+  30)  (1) 

(30)  +  H2O  +  S  -»  H2SO4 (2) 

2HNO3  +  S  -» 2NO  +  H2SO4  (3) 

2HN03  ( =  H2O,N205)  ->  2N02  +  H20  +  O     X  3  (4) 

(30)  +  H20  +  S  ->  H2S04  (5) 

6HNO3  +  S  -» 6NO2  +  2H2O  +  H2SO4  (6) 

The  reader  will  note  (cf.  p.  483)  that  a  separate  equation,  (3)  and  (6), 
must  be  made  for  the  formation  of  each  reduction  product.  If  NO 
and  N02  are  both  formed,  they  cannot  arise  from  the  same  molecule 
of  nitric  acid.  They  result  from  two  actions  which  are  independent, 
although  proceeding  concurrently  in  the  same  vessel  (cf.  p.  483). 
Thus  the  equation:  2HNO3  +  C  -»  H2O  +  CO2  +  NO •  +  NO2,  is 
a  misrepresentation.  It  implies  that  equimolar  quantities  of  the 
two  oxides  of  nitrogen  are  formed.  But  this  could  occur  only  by 
chance,  and  the  balance  would  be  destroyed  the  next  moment  by  the 
lowering  in  the  concentration  of  the  acid,  giving  the  advantage  to  the 
nitric  oxide. 

Oxidation  of  Compounds:  Aqua  Regia.  —  Compounds  like 
hydrogen  sulphide,  hydrogen  iodide,  and  sulphurous  acid,  which 
are  easily  oxidized,  interact  with  nitric  acid.  With  diluted  nitric 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  537 

acid  the  products  are  free  sulphur,  iodine,  and  sulphuric  acid,  re- 
spectively. 

The  mixture  of  nitric  acid  and  hydrochloric  acid  is  known  as 
aqua  regia.  Chlorine  is  set  free  by  the  oxidation  of  the  hydro- 
chloric acid, 

O 


Cl-jH  +  H  -  O  -JNJ  =  0  +  2H;C1->2H2O  +  C 

and  nitrosyl  chloride  NOC1  is  also  formed.  The  liquid  thus  con- 
tains several  oxidizing  agents,  nitric  acid,  hypochlorous  acid  (from 
C12  +  H20),  and  some  nitrous  acid.  It  is  frequently  used  in  analysis, 
for  example  to  oxidize  sulphur  (say,  in  cast  iron  or  in  minerals), 
the  sulphuric  acid  formed  being  estimated  by  precipitation  and 
weighing  of  barium  sulphate  (p.  440). 

Aqua  regia  (Lat.,  royal  water)  received  its  name  because  it  con- 
verted the  "noble"  metals,  gold  and  platinum,  into  soluble  com- 
pounds. This  it  does  because  the  free  chlorine,  in  presence  of  hydro- 
chloric acid  combines  to  form  the  exceedingly  stable  complex  ions 
(q.v.)  AuCl4~  (see  chlorauric  acid  HAuCl4)  and  PtCl6=,  the  negative 
ion  of  chloroplatinic  acid:  2HC1 4-  2C12  +  Pt -» H2PtCl6.  Since  the 
chloride-ion,  rather  than  the  free  chlorine,  is  concerned,  the  following 
equations  show  the  action  more  exactly: 

4H+      +N03~  ^  NO  +  2H2O  +  30x4  (1) 

Pt          -f-  40    ^Pf"-4-*-  X  3  (2) 

m pt+++++  6Cl-^PtCl6= X  3  (3) 

16H+  +  4NO3-+  3Pt  +  18Cr«*8HaO  +  4NO  +  3PtCl6= 


Mercuric  sulphide  HgS  is  attacked  because  mercuric  chloride  is  a 
very  slightly  ionized  substance,  particularly  in  presence  of  hydro- 
chloric acid,  with  which  it  combines. 

The  idea  that  a  nascent  form  of  chlorine  exists  here  is  even  more 
superfluous  than  is  the  similar  assumption  about  oxygen. 

NITROUS  ACID,  HYPONITROUS  ACID,  AND  THEIR  ANHYDRIDES 

Nitrites  and  Nitrous  Acid.  —  When  the  nitrates  of  potassium 
and  sodium  are  heated,  they  lose  one  unit  of  oxygen,  and  the  nitrites 
remain: 

2NaNO3  *+  2NaNO2  +  O2. 


538  INORGANIC  CHEMISTRY 

Commonly  lead  is  stirred  with  the  melted  nitrate  and  assists  in  the 
removal  of  the  oxygen.  The  litharge  PbO  which  is  formed  remains 
as  a  residue  when  the  sodium  nitrite  is  dissolved  for  recrystallization. 
When  an  acid  is  added  to  a  dilute  solution  of  a  nitrite,  a  pale-blue 
solution  containing  nitrous  acid  HN02  is  obtained.  The  acid  is  very 
unstable,  however,  and,  when  the  solution  is  warmed,  it  decomposes : 

3HN02  ->  HNO3  +  2NO  +  H2O. 

When  a  concentrated  solution  of  sodium  nitrite  is  acidified,  the  nitrous 
acid  decomposes  at  once,  and  a  brown  gas  containing  the  anhydride 
escapes: 

2H+  +  2NO2~  ^  2HNO2  *=»  H20  +  N2O3  T . 

This  behavior  distinguishes  a  nitrite  from  a  nitrate. 

Reducing  agents  deprive  nitrous  acid  of  part  or  all  of  its  oxygen: 

2H1  +  2HN02  (or  H20,N203)  ->  2H20  +  2NO  +  I2- 

Indigo  is  also  converted  by  it  into  isatin  (cf.  p.  476).  On  the  other 
hand,  oxidizing  agents  which  are  sufficiently  active,  like  acidified 
potassium  permanganate,  convert  nitrous  acid  into  nitric  acid: 

3H2SO4  +  2KMnO4  -»  K2SO4  +  2MnS04  +  3H20  (+  50)  (1) 

(50)  +  5HN02  ->  5HN03 (2), 

3H2SO4+2KMnO4+5HNO2^K2SO4+2MnSO4+3H20+5HNO3 
Nitrous  acid  is  much  used  in  the  making  of  organic  dyes. 

Nitrous  Anhydride  NZO3.  —  A  study  of  the  density  of  the  gas 
arising  from  the  decomposition  of  nitrous  acid  shows  that,  in  the 
gaseous  state,  the  anhydride  is  almost  entirely  dissociated: 

N203  <=>  NO  +  N02. 

When  the  mixture  is  led  through  a  U-tube  immersed  in  a  freezing 
mixture  at  —21°,  a  deep-blue  liquid  is  obtained  which  appears  to  be 
the  anhydride  itself.  This  begins  to  dissociate  before  reaching  its 
boiling-point,  and  at  +2°  gives  off  nitric  oxide. 

The  same  equimolar  mixture  of  the  two  gases  is  obtained  by  the 
action  of  water  on  nitrosylsulphuric  acid  (p.  433). 

Hyponitrous  Acid  H2]V2O2.  —  This  acid  is  formed  by  the  in- 
teraction of  hydroxylamine  and  nitrous  acid  in  aqueous  solution : 


OXIDES  AND  OXYGEN   ACIDS  OF  NITROGEN  539 

With  nitrate  of  silver,  the  yellow,  insoluble  silver  hyponitrite  Ag2N2C>2 
is  precipitated.  When  this  salt  is  shaken  with  an  ethereal  solution  of 
hydrogen  chloride,  the  acid  is  liberated,  and  the  insoluble  silver 
chloride  may  be  separated  by  filtration.  Finally,  evaporation  of  the 
ethereal  solution  leaves  hyponitrous  acid  as  a  white  solid.  It  explodes 
when  heated,  and  its  solution  in  water  is  an  exceedingly  feeble  acid. 
The  warm  aqueous  solution  decomposes  slowly,  giving  nitrous  oxide : 

H2N202  ->  H20  +  N20, 
and  this  change  is  not  capable  of  reversal. 

Nitrous  Oxide  N2O.  —  Nitrous  oxide  is  prepared  by  cautiously 
heating  ammonium  nitrate  (the  decomposition  is  exothermal),  or  a 
mixture  of  a  salt  of  ammonium  and  a  nitrate: 

NH4NO3  ->  2H20  +  N2O. 

The  steam  condenses,  and  the  nitrous  oxide  may  be  collected  over 
warm  water,  or  dried  and  compressed  into  steel  cylinders. 

Its  solubility  in  cold  water  is  considerable:  at  0°,  130  volumes  in 
100;  at  25°,  60  in  100.  In  dissolving,  the  gas  forms  no  compound 
with  water.  The  substance  melts  at  —102.3°,  and  boils  at  —89.8°. 
The  vapor  tension  of  the  liquid  at  0°  is  30.75  atmospheres;  at  12°, 
41.2  atmospheres;  and  at  20°,  49.4  atmospheres.  The  critical 
temperature  is  38.8°. 

A  glowing  splinter  of  wood  bursts  into  flame  in  nitrous  oxide, 
and  phosphorus,  sulphur,  and  other  combustibles,  burn  in  it  with 
much  the  same  vigor  as  in  oxygen.  In  all  cases  oxides  are  formed, 
and  nitrogen  is  set  free.  It  does  not  interact  with  nitric  oxide, 
however,  as  does  oxygen  (p.  529).  The  rapidity  with  which  bodies 
combine  with  oxygen  obtained  from  nitrous  oxide  is  doubtless  due  to 
the  fact  that  it  is  an  endothermal  compound,  and  the  heat  liberated 
by  its  decomposition  assists  the  ensuing  combustion: 

2N20  ->  2N2  +  02  +  2  X,  18,000  cal. 

It  is  to  be  noted  that  the  effect  of  the  heat  of  decomposition  will  be 
partly  offset  by  the  dilution  of  the  oxygen  with  nitrogen.  Yet  the 
proportion  of  nitrogen  to  oxygen  is  only  half  as  great  as  in  air,  so 
that  on  the  whole  the  conditions  are  much  more  favorable  to  combus- 
tion in  this  gas. 

Nitrous  oxide,  when  cold,  does  not  behave  like  free  oxygen. 
Metals  do  not  rust  in  it,  and  the  haemoglobin  of  the  blood  is  unable 


540  INORGANIC  CHEMISTRY 

to  use  it  as  a  source  of  oxygen.  It  was  Davy  who  first  observed  that 
nitrous  oxide  could  be  taken  into  the  lungs,  and  that,  since  it  furnished 
no  oxygen,  insensibility  followed  its  use.  By  suitable  admixture  of 
an  amount  of  air  sufficient  to  sustain  life,  it  is  employed  as  an  anaes- 
thetic for  minor  operations.  The  hysterical  symptoms  which 
accompany  its  use  caused  it  to  receive  the  name  of  " laughing  gas." 

Graphic  Formulae  of  Nitric  Acid  and  its  Derivatives:  Ex- 
plosives. —  The  following  equation  shows  the  graphic  formulae  of 
nitric  acid  and  of  ammonium  nitrate: 

5'^  *°  HX  ^° 

£>N-OH    +    H-O-Nj       M)     £>N-0-Nf       +    H2O 

H'  °  H'  ° 

The  structural  formula  of  the  latter  is  intended  to  explain  the  fact 
that  the  salt  is  able  to  exist  at  all,  by  representing  the  oxygen  and 
hydrogen  as  being  separated  from  one  another  and  attached  to  differ- 
ent nitrogen  units.  When  the  equilibrium  of  the  system  is  disturbed 
by  heating,  the  oxygen  and  hydrogen  unite  to  form  water,  an  arrange- 
ment which  is  much  more  stable,  and  nitrous  oxide  (p.  539)  escapes 
with  the  steam. 

The  behavior  of  nitroglycerine  and  guncotton  (p.  528),  as  well 
as  of  ammonium  nitrite  (p.  514),  is  explained  in  the  same  way.  These 
substances  are  made  by  actions  which,  like  the  above  neutralization, 
take  place  in  the  cold,  and  the  groups,  containing  the  oxygen  on 
the  one  hand  and  carbon  and  hydrogen  on  the  other,  become  quietly 
united  without  more  serious  interaction.  Thus  the  formation  of 
nitroglycerine  (p.  528)  appears  as  follows: 

H  H 

I  I 

H  -  C  -  OH  HO  -  N02  H  -  C  -  O  -  N02 

I  I 

H  -  C  -  OH    +    HO  -  NO2     ->     H  -  C  -  0  -  NO2    +    3H2O 

I  I 

H  _  C  -  OH  HO  -  N02  H  -  C  -  O  -  NO2 

I  I 

H  H 

When  the  nitroglycerine  is  heated,  or  receives  a  mechanical  shock, 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  541 

the  oxygen  all  unites  with  the  carbon  and  hydrogen,  and  the  nitrogen 
escapes: 

4C3H6(ON02)3  ->  12C02  +  10H20  +  6N2  +  02. 

That  nitroglycerine  is  a  more  sensitive  explosive  than  gunpowder 
is  due  to  the  fact  that,  in  the  former,  the  materials  required  for  the 
chemical  change  are  already  within  the  same  molecule,  whereas  in 
the  latter  (q.v.)  they  are  contained  in  the  separate  molecules  of  a  mix- 
ture. Even  after  the  most  careful  incorporation,  the  oxygen  of  the 
potassium  nitrate  can  hardly  be  uniformly  so  near  to  the  carbon, 
mechanically  mixed  with  the  salt,  as  are  these  elements  in  nitro- 
glycerine or  guncotton.  In  the  latter  the  oxidation  of  the  hydrogen 
and  carbon  is  intramolecular. 

•  Substances  like  hydrazoic  acid  (p.  521)  and  nitrogen  iodide  (p.  523) 
might  seem  to  constitute  a  third  kind  of  explosive.  Here  the  change 
consists  in  the  resolution  of  the  compound  into  its  constituents.  Still, 
if  we  consider  the  case  of  hydrazoic  acid,  for  example:  2N3H  — >  3N2 
+  H2,  we  see  that  the  action  consists,  after  all,  in  the  union  of  the  con- 
stituents to  form  the  more  stable  combinations  N2  and  H2.  It  is, 
therefore,  similar  in  principle  to  the  explosion  of  nitroglycerine. 

Smokeless  Powder  and  Dynamite.  —  Dried  guncotton  (p. 
528)  simply  burns  briskly  (deflagrates)  when  set  on  fire.  Whether 
wet  or  dry,  it  explodes,  but  only  from  a  suitable  shock,  such  as 
that  produced  by  fulminate  of  mercury  Hg(ONC)2,  used  in  per- 
cussion caps.  In  pure  form  it  is  used  only  in  torpedoes  or  sub- 
marine mines.  Like  nitroglycerine  (p.  541),  it  explodes  too  rapidly, 
and  would  burst  the  gun,  or  pulverize  the  ore  or  coal  if  used  for 
blasting.  Neither  of  these  substances  "explodes  downwards  only." 
The  explosion  strikes  the  air  with  equal  violence,  but  the  effect  on 
the  air  escapes  notice  because  it  is  not  permanent,  while  the  shattering 
of  a  rock  or  plate  of  steel  remains. 

Cordite,  one  variety  of  smokeless  powder,  is  made  by  dissolving 
guncotton  (65  parts),  nitroglycerine  (30  parts),  and  vaseline  (5 
parts)  in  acetone.  The  resulting  paste  is  rolled  out  and  cut  into 
small  pieces.  When  the  acetone  evaporates,  the  horny  cordite 
remains.  These  explosives  are  smokeless  because,  unlike  gun- 
powder and  T.N.T.,  they  yield  no  solids  when  they  decompose 
(see  equations). 

Various  forms  of  dynamite  are  made  like  cordite,  excepting  that 
sodium  or  ammonium  nitrate  and  sawdust  or  flour  are  added,  so 


542  INORGANIC  CHEMISTRY 

that  the  rate  of  explosion  may  be  regulated  and  the  coal  or  ore  may 
be  split  up,  but  not  shattered  or  pulverized. 

Plastics.  —  A  guncotton,  less  completely  "nitrated"  by  nitric 
acid,  when  worked  between  rollers  with  camphor  and  a  little  alcohol, 
gives  a  viscous  solution  (Parkes,  before  1855),  When  the  alcohol 
evaporates,  transparent,  colorless  celluloid  (first  made  by  Hyatt) 
remains.  The  moist  dough  is  rolled  into  sheets  to  make  photo- 
graphic films.  By  adding  dyes  and  "fillers,"  and  molding  the 
dough,  black  combs,  brush  handles,  white  knife  handles,  etc.,  can  be 
manufactured. 

Collodion  is  a  solution  of  the  same  guncotton  in  a  mixture  of 
alcohol  and  ether.  When  collodion  is  forced  through  minute  holes 
in  a  steel  dye,  the  threads  dry  as  they  come  out  and  can  be  wound 
on  spools.  Treatment  with  an  alkali  "denitrates"  the  threads, 
restoring  the  composition  to  that  of  the  original  cotton.  The  prod- 
uct, one  of  the  forms  of  artificial  silk,  is  at  least  as  brilliant  as  the 
real  article  (a  protein,  not  related  chemically  to  cotton),  and  sus- 
ceptible of  being  dyed  to  any  desired  tint. 

Balancing  Equations.  —  The  reader  should  practice  the  bal- 
ancing of  the  equations  for  oxidations  occurring  in  this  chapter, 
using  all  the  methods.  In  the  text,  we  have  used  the  anhydride 
plan  (p.  535)  and  that  of  partial  equations  (p.  538).  To  illustrate 
the  other  two  plans,  take,  for  example,  the  action  of  concentrated 
nitric  acid  on  copper  (p.  535). 

Positive  and  negative  valence  method  (p.  493).  Write  the  skeleton 
equation: 

Skeleton:         HNO3  +  Cu  -»  H2O  +  NO2  +  Cu(N03)2. 

We  perceive  that  on  the  left  the  valence  of  03  is  —6  and  of  H  is 
+  1:  that  of  N  is  therefore  +5.  That  of  Cu  is  zero.  On  the  right, 
the  valence  of  N  is  +4  and  of  Cu  +2.  Evidently,  2N  changing 
from  2X+5  to  2X+4  will  furnish  +2  for  the  copper.  We  note 
also  that  2N03  is  required,  without  change,  for  Cu(N03)2.  Hence, 
altogether  4HNO3  is  needed  on  the  left,  and  gives  2NO2 : 

Balanced:    4HN03  +  Cu  ->  2H2O  +  2N02  +  Cu(N03)2. 

Positive  electrical  charge  plan  (p.  496).  In  the  skeleton  equation 
(above)  we  first  separate  the  oxidizing  ions  and  their  products  from 
the  oxidized  substance  and  the  change  it  undergoes: 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  543 

NO3~  +  2H+  ->  N02°  -f  H20  +  ©    X  2 

Cu°  +  20  -»  Cu++ 

Cu°  +  2N03~  +  4H+  -5  2N02°  +  2H20  +  Cu++ 

The  first  partial  equation  produces  ©,  while  the  second  requires 
2©,  and  hence  the  former  is  multiplied  by  2  before  the  addition 
takes  place.  Since  NOa~  is  the  only  acid  radical  present,  it  is  under- 
stood that  cupric  nitrate  is  the  salt  formed. 

Active  ("  Nascent  ")  Hydrogen.  —  When  hydrogen  gas  is  led 
through  cold  nitric  acid,  little  or  no  action  occurs.  But  (p.  534) 
when  zinc,  which  interacts  with  acids  to  give  hydrogen,  is  placed 
in  nitric  acid  the  latter  is  reduced.  To  explain  the  apparently  greater 
activity  of  the  hydrogen  in  the  second  instance,  we  note  the  fact 
that  it  is  liberated  on  the  surface  of  the  zinc.  The  contact  (catalytic) 
effect  of  the  zinc  increases  its  activity.  Many  metals  have,  in  a 
greater  or  less  degree,  this  power  of  increasing  the  activity  of  hydro- 
gen. Thus,  hydrogen  absorbed  in  platinum  or  palladium  (p.  124), 
or  liberated  by  electrolysis  on  poles  made  of  these  metals,  reduces 
nitric  acid  readily.  Other  elements,  such  as  the  oxygen  in  making 
sulphur  trioxide  (p.  428),  are  also  rendered  more  active  by  contact 
agents  like  platinum. 

This  more  active  state  of  hydrogen  is  described  as  the  nascent 
state,  because  it  happens  to  be  a  common  condition  of  hydrogen 
when  associated  with  substances  which  produce  it.  The  active 
state  has,  however,  no  necessary  connection  with  such  an  immediately 
preceding  act  of  liberation,  as  the  platinum  and  sulphur  trioxide 
illustrations,  and  the  following  experiment  [Lect.  exp.]  show:  Three 
test-tubes  are  filled  with  very  dilute  potassium  permanganate  solu- 
tion. Zinc  dust,  added  to  one,  generates  hydrogen  and  causes 
decolorization.  A  little  platinum  black  is  added  to  the  second,  and 
hydrogen  gas  is  led  through  this  and  the  third.  The  contact  action 
of  the  platinum  enables  the  hydrogen  quickly  to  reduce  the  per- 
manganate, while  the  third  portion  remains  unaltered. 

The  term  nascent  hydrogen  is  used  in  different  senses,  in  a  very  confusing 
way.  (1)  It  may  mean  nascent,  literally,  that  is,  newly  born  or  liberated.  (2) 
It  is  used  also  to  mean  different-from-ordinary,  or,  in  fact,  an  allotropic  form  of 
hydrogen.  (3)  It  is  often  limited  to  mean  one  particular  allotrope,  namely, 
atomic  hydrogen.  (4)  It  is  used  by  Haber  and  others,  as  we  have  used  it  above,  to 
mean  hydrogen  activated  by  contact  with  a  metal.  (5)  Finally,  its  activity  is 
explained  as  being  due  to  the  larger  amount  of  free  energy  contained  in  zinc  plus 


544  INORGANIC  CHEMISTRY 

acid  plus  reducing  agent,  as  compared  with  the  free  energy  contained  in  free 
hydrogen  plus  reducing  agent.  The  last  is  identical  with  the  explanation  of  the 
activity  of  oxidizing  agents  (p.  479).  The  word  nascent  is,  of  course,  a  misnomer, 
excepting  in  connection  with  (1). 

The  following  statements  will  enable  us  to  determine  which  of  these  five 
conceptions  are  most  in  accord  with  experimental  facts.  This  form  of  hydrogen 
has  never  been  observed  or  isolated  as  a  substance  (against  1,  2,  3).  If  it  is  an 
allotropic  form  (2,  3),  its  degree  of  activity  can  be  defined  quantitatively.  In 
point  of  fact,  however,  concentrated  sulphuric  acid  gives,  with  copper,  sulphur 
dioxide  (p.  425),  and  with  zinc,  hydrogen  sulphide  (p.  119),  so  that  the  hydrogen 
(if  it  is  the  active  agent)  is  much  more  active  in  the  second  case.  Again,  in  elec- 
trolyzing  a  dilute  acid,  the  dissolved  atmospheric  oxygen  is  reduced  to  hydrogen 
peroxide,  provided  the  electrode  is  made  of  platinum  (p.  317),  but  not  if  it  is  made 
of  carbon.  We  saw  also  that  hydrogen  liberated  by  electrolysis  on  a  surface  of 
mercury  gives  a  better  yield  of  hydroxylamine  (p.  522),  than  will  any  other 
metallic  electrode.  Yet  the  hydrogen  must  be  the  same  in  all  cases  (against  1, 
2,  3;  favors  4).  Again,  hydrochloric  acid  and  some  nitric  acid,  with  zinc,  give 
ammonia  (p.  534),  with  magnesium  no  ammonia,  with  tin  ammonia  and  hydroxyl- 
amine (p.  522).  Here  again  the  hydrogen  is  the  same,  but  the  metallic  contact 
agent  is  different,  and  the  free  energy  of  the  acid  with  each  metal  is  different 
(against  1,  2,  3;  favors  4  and  5).  Some  one  found  that  nitrous  oxide  N2O  could 
be  prepared  by  heating  dry  potassium  nitrate  with  anhydrous  formic  acid  HCC^H : 
2KNO3  +  6HCO2H  ->  N2O  t  +  4CO2  +  5H2O  +  2KCO2H.  Formic  acid,  as  a 
whole,  has  undoubted  reducing  power,  so  why  drag  in  nascent  hydrogen,  as  the 
discoverer  did  in  this  case?  Finally,  since  hydrogen  and  chlorine  do  not  unite  in 
the  cold,  when  sulphuric  acid  and  common  salt  give  hydrogen  chloride,  to  be 
consistent  we  must  suppose  that  nascent  hydrogen  and  nascent  chlorine  were 
formed  and  combine.  In  other  words,  every  union  of  two  elements,  other  than 
direct  union,  must  be  explained  by  nascent  action,  although  in  double  decomposi- 
tion this  logical  necessity  is  uniformly  overlooked. 

There  seems  to  be  no  question  that  contact  with  different  metals  confers 
on  free  hydrogen  the  ability  to  produce  different  chemical  changes  in  the  same 
substance.  It  is  also  clear  that  the  extra  energy  with  which  hydrogen  is  delivered 
from  some  chemical  actions,  as  compared  with  others,  must  appear  to  give  it 
different  degrees  of  activity. 

The  Principle  of  Transformation  by  Steps.  —  It  may  have 
occurred  to  the  reader  as  strange  that  it  should  be  possible  to  make 
nitric  anhydride  by  distilling  a  warm  mixture  (p.  527)  when  the  prod 
uct  decomposes  spontaneously,  even  when  kept  in  the  cold.  How 
can  a  compound  be  fitted  together  under  certain  conditions,  when 
under  the  same  or,  even,  under  more  favorable  conditions  it  proves 
to  be  incapable  of  continued  existence?  We  should  expect  rather 
that  obtaining  the  products  of  its  decomposition  would  have  been  the 
only  result  of  the  effort  to  make  it. 


OXIDES  AND  OXYGEN  ACIDS  OF  NITROGEN  645 

Extraordinary  as  this  fact  appears  to  be,  it  is  nevertheless  very 
commonly  encountered.  Perchloric  acid  is  made  by  a  distillation 
(p.  484)  and  afterwards  breaks  up  of  its  own  accord.  So,  also,  hypo- 
chlorites  are  formed  first  and  can  be  isolated.  But  under  the  same 
conditions  the  further  transformation  to  chlorates  will  occur  (p.  480). 
A  simple  case  is  that  of  sulphur  made  by  precipitation  (p.  422)  at 
the  ordinary  temperature.  Although  it  is  naturally  solid  below  119°, 
yet,  when  first  thrown  down,  it  is  in  the  form  of  liquid  droplets  which, 
if  undisturbed,  may  remain  fluid  for  weeks.  Similarly,  sulphur  vapor 
condenses  on  glass  in  drops  which  remain  liquid  for  years,  unless  they 
are  touched  or  rubbed.  Finally,  a  supersaturated  solution  (p.  193) 
is  not  unlike  cold  liquid  sulphur. 

In  all  these  cases  there  is  a  possibility  of  further  change,  which, 
when  it  comes,  will  liberate  heat  or  some  other  form  of  energy.  Thus, 
heat  is  set  free  when  the  liquid  sulphur  is  precipitated.  The  amount 
of  heat  would  have  been  greater,  by  the  heat  of  fusion  of  the  sulphur, 
if  solidification  had  occurred  simultaneously.  But,  in  spite  of  the 
existence  of  this  justification  for  the  final  step,  this  step  is  not  taken. 
So,  the  decomposition  of  the  vapor  of  the  perchloric  acid  or  of  the 
nitric  anhydride  would  have  added  to  the  amount  of  energy  liberated 
as  heat,  but  this  additional  step  was  postponed.  In  other  words, 
transformations  which,  proceed  spontaneously  and  with  evolution  of 
heat  may  go  forward  by  steps,  when  there  are  intermediate  sub- 
stances, or  allotropic  forms,  capable  of  existence.  This  is  known  as 
the  principle  of  transformation  by  steps,  and  was  first  formulated  by 
Ostwald. 

Exercises.  —  1.  Make  the  equation  for  the  interaction  of  fer- 
rous chloride,  hydrochloric  acid,  and  nitric  acid  (p.  529),  and  for  all 
the  actions  concerned  when  the  test  for  a  nitrate  (p.  530)  is  applied 
to  sodium  nitrate.  What  volume  (at  0°  and  760  mm.)  of  NO  is 
obtained  from  one  formula-weight  of  nitric  acid  (p.  265,  Ex.  4)? 

2.  Should  you  classify  as  molecular  compounds  (p.  530) :  Chlorine 
hydrate,  ammonium  hydroxide,  KIs  (p.  276),  sulphurous  acid,  sodium 
tetrasulphide  (p.  422)?    Justify  your  answer. 

3.  At  70°,  what  proportions  of  the  molecules  of  nitrogen  tetroxide 
are  in  the  forms  of  N02  and  N2O4  respectively  (p.  532)?    At  the  same 
temperature  what  fraction  of  the  material,  by  weight,  is  in  the 
former  condition?    What  are  the  relative  volumes  of  the  tetroxide, 
and  of  the  nitric  oxide  and  oxygen  obtained  by  its  decomposition 
(p.  532)? 


546  INORGANIC  CHEMISTRY 

4.  Make  an  equation  showing  the  production  of  nitrous  oxide  by 
the  action  of  zinc  on  nitric  acid. 

5.  Make  the  correct  equations  showing  the  formation  of  nitric 
oxide  and  nitrogen  tetroxide  by  the  interaction  of  carbon  and  nitric 
acid  (p.  536). 

6.  Justify  the  graphic  formula  assigned  to  nitric  acid  (p.  540). 

7.  Using  the  anhydride  method  (p.  496),  make  the  equations  for 
the  interactions  of  N2O4  and  water  (p.  532)  and  of  permanganic 
acid  and  nitrous  acid  (p.  538). 

8.  In  the  action  of  zinc  on  dilute  nitric  acid  (p.  534),  why  is  not 
the  ammonia  given  off  as  a  gas?    How  should  you  show  that  it  was 
formed  at  all? 

9.  Make  equations  for  the  interaction  of  iron  with  diluted  and 
with  concentrated  nitric  acid,  respectively  (p.  535).     The  iron  gives 
ferric  nitrate  Fe(NO3)3 

10.  Give  the  three  ways  in  which  nitrates  decompose  when 
heated,  with  one  equation  illustrating  each. 

11.  Make  all  the  equations  for  oxidations  on  pp.  535  and  536, 
using  the  methods  illustrated  on  p.  542, 


CHAPTER  XXVII 
PHOSPHORUS 

The  Chemical  Relations  of  the  Element.  —  There  are  many 
things  in  the  chemistry  of  phosphorus  and  its  compounds  which 
remind  us  of  nitrogen.  Yet  these  are  largely  referable  to  the  fact 
that  the  elements  are  both  non-metals  and  both  have  the  same 
valences,  viz.,  three  and  five.  The  behavior  of  the  compounds  is  often 
very  different.  For  the  present  it  is  sufficient  to  say  that  both  give 
compounds  with  hydrogen,  NH3  and  PH3,  and  both  yield  oxides  of 
the  forms  X2O3,  X2O4,  and  X2O5.  The  first  and  last  of  these  oxides 
are  acid-forming,  and  phosphorus,  therefore,  gives  acids  corresponding 
to  nitrous  and  nitric  acids,  although  there  is  more  variety  in  the 
proportion  of  water  combined  with  the  anhydride  (cf.  p.  488).  The 
element  is  thus  non-metallic  (see  Comparison  with  nitrogen  and  with 
sulphur,  p.  564). 

Occurrence.  —  This  element  is  found  in  nature  in  the  form  of 
phosphates.  Calcium  phosphate  Ca3(PO4)2  forms  26  per  cent  of 
the  bones  and  teeth,  and  it  occurs  in  all  fertile  soils.  It  consti- 
tutes a  large  part  of  the  "phosphate  rock"  of  Georgia,  Florida, 
the  Carolinas,  Tennessee,  and  of  Algeria  and  Tunis.  A  conspicuous 
mineral  related  to  this  substance,  apatite  (Gk.  aTraraw,  I  cheat), 
CasF(PO4)3  and  Ca5Cl(PO4)3,  is  found  in  large  quantities  in  Canada, 
and  is  a  component  of  many  rocks.  Complex  organic  compounds  of 
phosphorus,  such  as  lecithin  and  some  proteins,  are  essential  con- 
stituents of  the  muscles,  the  nerves,  and  the  brain.  Amongst  foods, 
egg-yolks  and  beans  contain  a  large  proportion. 

Preparation.  —  Brand,  merchant  and  alchemist,  of  Hamburg, 
discovered  phosphorus  (1669)  by  distilling  the  residue  from  evapo- 
rated urine,  in  the  course  of  his  search  for  the  Philosopher's  stone. 
The  mode  of  preparing  it  from  bone-ash  was  first  published  by  Scheele 
(1771).  Green  bones,  after  the  gelatine  has  been  extracted  from  them, 
by  means  of  water  boiling  under  pressure,  are  subjected  to  destructive 
distillation,  a  process  which* yields  bone-oil.  The  residue  is  a  mixture 

M7 


548 


INORGANIC   CHEMISTRY 


of  carbon  (q.v.)  and  calcium  phosphate.  It  is  used  by  sugar  refiners 
as  a  decolorizer.  When  its  powers  in  this  direction  have  been  ex- 
hausted, it  is  calcined  —  that  is  to  say,  all  the  combustible  matter  is 
burned  out  of  it,  —  and  the  product  is  bone-ash,  which  contains 
about  83  per  cent  of  calcium  phosphate.  Formerly  this  was  used  in 
making  phosphorus,  but  now  the  less  expensive  calcium  phosphate 
of  fossil  origin  is  employed. 

A  mixture  of  powdered  bone-ash  or  calcium  phosphate  and  sulphu- 
ric acid  (sp.  gr.  1.5  to  1.6)  is  heated  with  steam  and  stirred  in  a 
wooden  vat:  Ca3(PO4)2  +  3H2SO4 ->  2H3PO4  +  3CaSO4  j.  The  cal- 
cium sulphate  is  partly  precipitated  during  the  heating.  The  liquid 
obtained  by  filtration  is  evaporated  in  leaden  pans.  During  this 
process  most  of  the  remainder  of  the  calcium  sulphate  is  deposited  and 
a  syrupy,  crude  phosphoric  acid  is  obtained.  This  acid  is  mixed  with 
sawdust,  or  carbon  in  some  form,  and  the  mixture  is  first  dried 
and  then  distilled  in  earthenware  retorts.  The  phosphoric  acid 
first  loses  water  and  turns  into  metaphosphoric  acid,  then  the  latter 
is  reduced  by  the  carbon,  carbon  monoxide 
and  phosphorus  vapor  passing  off: 

H3P04  -»  H2O  +  HPO3. 
2HP03  +  6C  -» H2  +  6COT  +  2P|. 

A  white  heat  is  required  for  the  distillation,  and 
a  pipe  from  the  tubular  clay  retort  conducts  the 
vapors  into  cold  water,  in  which  the  phosphorus 
collects. 

A  much  simpler  and  more  modern  process 
depends  on  the  use  of  the  electric  furnace 
(Fig.  125).  The  calcium  phosphate  is  mixed 
with  the  proper  proportions  of  carbon  and 
silicon  dioxide  (sand),  and  the  mixture  is  introduced  continuously 
into  the  furnace.  The  discharge  of  an  alternating  current  between 
carbon  poles  produces  the  very  high  temperature  which  the  action 
requires.  The  calcium  silicate  which  is  formed  fuses  to  a  slag,  and 
can  be  withdrawn  at  intervals.  The  gaseous  products  pass  off 
through  a  pipe  and  the  phosphorus  is  caught  under  water: 

Ca3(P04)2  +  3SiO2  +  5C  -»  3CaSi03  +  5COt  +  2P  T-. 

We  may  regard  the  phosphate  as  being  composed  of  two  oxides,  3CaO, 
It  thus  appears  that  the  calcium  oxide  has  united  with  the 


FIG.  125. 


PHOSPHORUS  549 

silica,  which  is  an  acid  anhydride  (cf.  p.  431) :  CaO  +  SiO2  — >  CaSi03, 
while  the  phosphoric  anhydride  has  been  reduced. 

The  phosphorus,  after  purification,  is  cast  into  sticks  in  tubes  of 
tin  or  glass,  standing  in  cold  water. 

The  Electric  Furnace.  —  By  an  electric  furnace  is  understood 
an  electro-thermal  arrangement  in  which  the  heat  produced  by  some 
resistance  offered  to  the  current,  such  as  that  of  an  air-gap  between 
the  carbons,  is  used  to  produce  chemical  change.  Electrolysis  plays 
no  part  in  the  phenomena,  and  an  alternating  current,  which  -can  pro- 
duce no  electrolytic  decomposition,  is  generally  employed.  The 
restricted  area  within  which  the  heat  is  developed  makes  possible  the 
attainment  of  a  high  temperature  (see  Calcium  carbide). 

Physical  Properties.  —  There  are  at  least  two  allotropic  forms 
(p.  315)  of  phosphorus,  known  as  white  phosphorus  and  red  phos- 
phorus. White  phosphorus,  prepared  as  described  above,  is  at  first 
transparent  and  colorless,  but  after  exposure  to  light  acquires  a 
superficial  coating  of  the  red  variety.  Its  density  is  1.83.  It  melts 
at  44°,  and  boils  at  287°.  Its  molecular  weight  at  313°  is  128,  and 
at  a  red  heat  119.8.  As  the  atomic  weight  is  31,  the  formula, 
within  this  range,  is  P4.  At  1700°  the  value  91.2  indicates  a  partial 
dissociation  into  ?2.  In  solution  the  formula  is  P4.  White  phos- 
phorus is  very  soluble  in  carbon  disulphide,  less  soluble  in  ether 
and  other  organic  solvents,  and  insoluble  in  water.  It  is  exceedingly 
poisonous,  less  than  0.15  g.  being  a  fatal  dose,  and  is  an  ingredient  in 
roach  paste  and  rat  poison  (a  mixture  with  lard  as  solvent,  and  flour). 
Continued  exposure  to  its  vapor  causes  necrosis,  a  disease  from 
which  matchmakers  are  liable  to  suffer.  The  jawbones  and  teeth 
are  particularly  liable  to  attack. 

Red  phosphorus  is  a  red  powder  consisting  in  part  of  small  tabular 
crystals.  It  is  obtained  by  heating  white  phosphorus  to  about  250° 
in  a  vessel  from  which  air  is  excluded.  The  change  is  much  more 
rapid  at  slightly  higher  temperatures.  Since  a  great  amount  of  heat 
is  evolved  in  the  transformation,  the  phosphorus  is  closely  confined 
to  prevent  volatization.  A  trace  of  iodine  accelerates  the  trans- 
formation, and  it  then  takes  place  even  in  the  cold. 

Red  phosphorus  does  not  melt,  but  passes  directly  into  vapor.  Its 
vapor  is  identical  with  that  of  white  phosphorus.  It  is  insoluble  in 
carbon  disulphide  and  other  solvents.  It  is  not  poisonous,  and, 
unlike  white  phosphorus,  does  not  require  to  be  kept  under  water  to 


550  INORGANIC  CHEMISTRY 

avoid  spontaneous  combustion.  Its  melting-point,  in  a  closed 
vessel,  varies,  being  550°  when  heated  very  slowly,  and  600° 
when  heated  rapidly.  This  shows  that  it  is  a  solid  solution,  prob- 
ably of  the  white  variety  in  a  less  active  kind.  Hence,  its  proper- 
ties are  variable,  e.g.,  density  from  2.05  to  2.34.  Bridgeman,  by 
heating  white  phosphorus  at  200°  under  a  pressure  of  1200  kg./cm.,2 
has  obtained  a  black,  lustrous  phosphorus  (sp.  gr.  2.69),  which  is  a 
third  variety,  and  has  a  fair  electrical  conductivity. 

Chemical  Properties  of  White  Phosphorus.  —  White  phos- 
phorus unites  directly  with  the  halogens  with  great  vigor.  It  unites 
slowly  with  oxygen  in  the  cold,  and  with  sulphur  and  many  metals 
when  the  materials  are  heated  together.  The  slow  union  of  cold 
phosphorus  with  atmospheric  oxygen  is  accompanied  by  the  evolution 
of  light,  although  the  temperature  is  not  such  as  we  usually  associate 
with  incandescence.  Hence  the  word  phosphorescence.  The  name 
of  the  element  (Gk.  <££>s,  light;  <l>ep<»,  I  bear)  records  this  property. 
Apparently  the  chemical  energy,  transformed  in  connection  with  the 
oxidation,  is  converted,  in  part  at  least,  into  radiant  energy  instead  of 
completely  into  heat.*  A  curious  fact  in  connection  with  the  lumi- 
nosity and  concomitant  oxidation  of  phosphorus  is  that  these  occur- 
rences depend  upon  the  concentration  of  the  oxygen  gas  as  well  as 
upon  the  temperature.  Thus,  phosphorus  does  not  shine  or  oxidize 
in  pure  oxygen  below  27°.  If  the  concentration  of  the  oxygen 
is  reduced  to  200  mm.  or  less  by  means  of  a  pump,  or  by  mixing 
with  an  indifferent  gas  such  as  nitrogen,  phosphorescence  becomes 
perceptible  at  the  ordinary  temperature.  This  explains  the  lumi- 
nosity shown  in  the  air.  At  lower  temperatures,  lower  pressures 
have  to  be  used.  The  phosphorescence  may  be  destroyed  by  the 
vapor  of  turpentine  and  other  substances.  All  these  phenomena 
are  probably  due  to  the  intermediate  formation  of  phosphorus 
trioxide,  the  vapor  of  which  shows  the  same  effects.  The  slow 
oxidation  of  phosphorus  is  accompanied  by  the  production  of  ozone, 
but  the  nature  of  the  action  is  still  unknown  (cf.  p.  311). 

*  The  same  production  of  light  from  chemical  action  in  a  cold  body  is  seen  in 
the  luminosity  of  certain  parts  of  fireflies  and  some  species  of  fish.  In  many 
violent  chemical  changes  the  light  given  out  is  conspicuously  more  intense  than 
that  proper  to  the  temperature  produced  (cf.  p.  94),  and  must  come,  therefore, 
in  part,  directly  from  the  chemical  energy.  Thus,  burning  magnesium  has  a 
temperature  of  about  1350°,  while  the  production  of  light  of  the  same  character., 
by  mere  incandescence,  would  require  a  temperature  of  about  5000°. 


PHOSPHORUS  551 

The  difference  in  behavior  of  pure  and  diluted  oxygen  may  be  shown  by  pour- 
ing a  solution  of  phosphorus  in  carbon  disulphide  on  to  two  strips  of  filter  paper. 
One  of  the  strips,  hung  in  the  air,  catches  fire  as  soon  as  the  evaporation  of  the 
solvent  has  exposed  a  large  area  of  finely  divided  phosphorus.  The  other,  hung 
in  a  jar  of  oxygen,  remains  unaffected,  but  becomes  ignited  instantly  upon  re- 
moval from  the  jar  [Lect.  exp.]. 

Chemical  Properties  of  Red  Phosphorus.  —  This  variety  of 
the  element,  since  it  is  formed  with  evolution  of  heat,  contains  less 
energy  than  does  white  phosphorus  and  is  much  less  active.  It  does 
not  catch  fire  in  the  air  below  240°,  while  ordinary  phosphorus  ignites 
at  35-45°.  Indeed,  it  is  the  vapor  that  begins  to  combine  with 
oxygen,  and  this  behavior  is  only  an  independent  proof  of  the  low 
vapor  tension  of  the  red  variety.  When  the  vapor  tension  of  white 
phosphorus  has  reached  760  mm.  (at  287°,  the  b.-p.),  that  of  red 
phosphorus  is  almost  imperceptible. 

Red  phosphorus  is  to  be  regarded  as  the  normal,  stable  form  of 
phosphorus.  The  fact  that  yellow  phosphorus  can  be  kept  a  great 
length  of  time,  and  is  changed  but  slowly  on  exposure  to  light,  only 
shows  that  the  transformation  into  a  stabler  condition  is  retarded  by 
the  lowness  of  the  temperature  (cf.  p.  126).  The  relation  between  the 
two  varieties  of  phosphorus  is  quite  distinct  from  that  between  rhom- 
bic and  monoclinic  sulphur  (p.  411).  In  the  latter  case  there  is  a 
definite  temperature  of  transformation  (96°)  above  which  one  form 
completely  disappears,  and  below  which  the  other  form  is  incapable  of 
permanent  existence.  With  the  varieties  of  phosphorus  no  such  point 
of  transformation  exists,  because  with  phosphorus  the  two  forms 
are  miscible,  while  with  sulphur  they  are  not.  It  is  only  by  con- 
densing the  vapor  that  the  yellow  kind  is  obtained.  The  production 
from  the  vapor  (also  P4)  of  the  white  solid  at  low  temperatures, 
instead  of  the  red  solid  whose  formation  would  be  accompanied  by  a 
larger  liberation  of  heat,  is  simply  an  illustration  of  the  principle  of 
transformation  by  steps  (p.  544).  At  1200°,  however,  the  vapor 
deposits  red  phosphorus. 

The  term  allotropic  modifications  is  applied  to  oxygen  and  ozone  (q.v.)  which 
are  certainly,  and  to  red  and  white  phosphorus  which  are,  probably,  chemically 
distinct  substances.  It  is  used  also  of  rhombic  and  monoclinic  sulphur,  where 
the  difference  is  purely  physical.  But  it  is  applied  only  to  simple  substances, 
although  many  compounds  show  several  forms  (cf.  Ammonium  nitrate),  exactly 
as  does  sulphur.  In  short,  it  has  at  present  no  scientific  value,  for  it  covers  a 
heterogeneous  mass  of  phenomena  which,  in  part,  still  await  elucidation.  If 
allotropic  modifications  were  to  be  denned  as  substances  (p.  4)  composed 


552  INORGANIC  CHEMISTRY 

of  the  same  materials,  but  possessing  different  proportions  of  free  or 
available  energy,  and,  therefore,  different  physical  properties  and  different 
degrees  of  chemical  activity,  which  is  apparently  the  sense  in  which  the  ex- 
pression is  commonly  employed,  then  ice,  water,  and  steam  would  be  examples 
of  such  substances.  In  each  of  the  above  four  illustrations,  the  second  (in  the 
case  of  water,  the  third)  is  the  more  active  form. 

Uses  of  Phosphorus,  Matches.  —  The  greater  part  of  the 
phosphorus  of  commerce  is  employed  in  the  manufacture  of  matches. 
The  first  articles  of  this  sort  (1812)  were  sticks  coated  with  sulphur 
and  tipped  with  a  mixture  of  potassium  chlorate  and  sugar.  For 
ignition  they  were  dipped  into  a  bottle  containing  asbestos  moistened 
with  concentrated  sulphuric  acid.  Matches  involving  the  use  of 
phosphorus  (1827)  have  now  displaced  all  others.  In  making  common 
matches  which  strike  on  any  rough  surface,  the  sticks  are  first  dipped 
in  melted  sulphur  or  paraffin  to  the  extent  of  about  half  an  inch.  The 
head  is  often  composed  of  manganese  dioxide  or  lead  dioxide  Pb02, 
and  a  little  potassium  chlorate,  which  supply  oxygen,  a  small  propor- 
tion of  white  phosphorus,  or  of  a  sulphide  of  phosphorus  P4S3  which  is 
readily  ignited  by  friction,  and  antimony  trisulphide  (combustible) 
with  dextrine  or  glue.  A  paste  made  of  these  materials  is  spread 
evenly  upon  a  slab,  and  the  prepared  sticks  fixed  in  a  frame  are  dipped 
once  or  twice  in  the  mixture.  On  account  of  the  danger  of  necrosis 
amongst  the  workers,  the  use  of  white  phosphorus  is  forbidden  by 
law  in  Sweden,  France,  Great  Britain,  and  Switzerland,  and  is  pre- 
vented by  a  tax  of  two  cents  per  100  matches  in  the  United  States. 

In  the  case  of  "safety"  matches,  the  mixture  upon  the  head  is  not 
easily  ignited  by  itself.  It  is  composed  of  potassium  chlorate  or  di- 
chromate,  some  sulphur  or  antimony  trisulphide,  and  a  little  powdered 
glass  to  increase  the  friction,  all  held  together  with  glue.  Upon  the 
rubbing  surface  on  the  box  is  a  thin  layer  of  antimony  trisulphide 
mixed  with  red  phosphorus  and  glue.  The  friction  converts  a  little 
of  the  red  phosphorus  into  vapor,  which  catches  fire  readily.  To 
prevent  smoldering  of  the  burned  matches,  the  upper  ends  of  the 
sticks  are  sometimes  soaked  in  a  solution  of  alum  or  sodium  phosphate. 

Phosphine.  —  Three  hydrides  of  phosphorus  are  known.  These 
are,  phosphine  PHs  (a  gas),  a  liquid  hydride  P2H4,  which  is  pre- 
sumably the  analogue  of  hydrazine  N2H4,  and  a  solid  hydride  P4H2. 

Phosphine  PH3  does  not  seem  to  be  produced  under  ordinary 
circumstances  by  the  direct  union  of  the  elements.  It  is  formed 


PHOSPHORUS  553 

slowly,  however,  with  active  hydrogen,  from  zinc  and  hydrochloric 
acid  at  70°,  and  white  phosphorus.  The  gas  may  be  made  by  boiling 
white  phosphorus  with  potassium  hydroxide  solution  in  an  apparatus 
similar  to  that  used  for  generating  hydrogen.  Potassium  hypo- 
phosphite  is  formed  at  the  same  time: 

3KOH  +  4P  +  3H20  ->  3KH2P02  +  PH3  ?• 

The  gas  made  in  this  fashion  contains  a  little  of  the  vapor  of  the 
liquid  hydride,  which  is  spontaneously  inflammable,  and  consequently 
the  bubbles  of  the  mixture  catch  fire  when  they  reach  the  surface  of 
water  in  the  trough :  PH3  +  2O2  — >  H3P04.  In  still,  moist  air, 
the  fog  of  droplets  of  phosphoric  acid  solution  form  smoke  rings. 
To  avoid  explosions,  the  air  in  the  flask  must  be  displaced  by  hydrogen 
or  illurninating-gas  before  heat  is  applied.  This  product  contains 
also  free  hydrogen,  in  increasing  quantities  as  the  action  goes  on,  in 
consequence  of  the  reduction  of  the  water  and  potassium  hydrox- 
ide by  the  potassium  hypophosphite:  KH2P02  +  KOH  -{-  H2O  — > 
K2HP(>4  +  2H2.  Potassium  phosphate  is  formed. 

The  simplest  method  of  preparing  the  gas  is  by  the  action  of  water 
upon  calcium  phosphide: 

Ca3P2  +  6H20  ->  3Ca(OH)2  +  2PH3. 

This  action  is  analogous  to  that  of  water  upon  magnesium  nitride 
(p.  514),  by  which  ammonia  is  produced.  In  consequence  of  the  fact 
that  calcium  phosphide  is  a  substance  of  irregular  composition,  a  mix- 
ture of  all  three  hydrides  is  generally  obtained.  By  passing  the  gas 
through  a  strongly  cooled  delivery  tube,  however,  the  liquid  com- 
pound is  condensed  and  fairly  pure  phosphine  passes  on. 

Phosphine  is  a  colorless  gas,  which  is  easily  decomposed  by  heat 
into  its  elements.  It  is  exceedingly  poisonous  and,  unlike  ammonia, 
it  is  insoluble  in  water,  and  produces  no  basic  compound  corresponding 
to  ammonium  hydroxide  when  brought  in  contact  with  this  sub- 
stance. It  resembles  ammonia,  formally  at  least,  in  uniting  with  the 
hydrogen  halides  (see  below).  It  differs  from  ammonia,  however, 
inasmuch  as  it  does  not  unite  with  the  oxygen  acids.  Phosphine 
acts  upon  solutions  of  some  salts,  precipitating  phosphides  of  the 
metals: 

3CuSO4  +  2PH3  ->  Cu3P2  +  3H2S04. 

The  liquid  hydrogen  phosphide  boils  at  57°.  The  molecular 
weight,  as  determined  by  the  density  of  its  vapor,  shows  the  formula  to 


554  INORGANIC  CHEMISTRY 

be  P2H4.  It  forms  no  salts,  and  is  therefore  quite  unlike  hydrazine. 
When  exposed  to  light,  it  decomposes,  giving  phosphine  and  the 
solid  hydride. 

Phosphonium  Compounds.  —  Hydrogen  iodide  unites  with 
phosphine  to  form  a  colorless  solid,  crystallizing  in  beautiful,  highly 
refracting,  square  prisms:  PH3  +  HI  — •>  PH4I.  Hydrogen  chloride 
combines  similarly  with  phosphine,  but  only  when  the  gases  are  cooled 
by  a  freezing  mixture,  or  are  brought  together  under  a  total  pressure 
of  18  atmospheres  at  14°.  When  the  pressure  is  released,  rapid 
dissociation  occurs.  This  dissociation  is  one  of  the  many  cases  where 
an  action  which  absorbs  heat,  nevertheless  goes  on  spontaneously 
(c/.  p.  35).  The  indispensable  fall  in  the  energy  of  the  system  takes 
place  by  virtue  of  the  expansion  of  the  constituents,  and  in  amount 
this  more  than  offsets  the  heat  acquired. 

In  imitation  of  the  ammonia  nomenclature,  these  substances  are 
called  phosphonium  iodide  and  phosphonium  chloride  PH4C1.  They 
are  entirely  different,  however,  from  the  corresponding  ammonium 
derivatives,  for  the  PH4+  ion  is  unstable.  When  brought  in  contact 
with  water,  they  decompose  into  their  constituents,  the  hydrogen 
halide  going  into  solution,  and  the  phosphine  being  liberated  as 
a  gas. 

Halides  of  Phosphorus.  —  The  existence  of  the  following 
halides  has  been  proved  conclusively: 

......  P2I4  (solid) 

PF3  (gas)        PC13  (liquid)          PBr3  (liquid)        PI3  (solid) 
PF5  (gas)        PC15  (solid)  PBr5  (solid)  

These  substances  may  all  be  formed  by  direct  union  of  the  elements. 
They  are  incomparably  more  stable  than  are  the  similar  compounds  of 
nitrogen.  They  are  all  completely  hydrolyzed  by  water,  and  each 
gives  an  oxygen  acid  of  phosphorus  and  the  hydrogen  halide  (see 
below).  This  action  was  used  in  the  preparation  of  hydrogen  bro- 
mide (p.  272)  and  hydrogen  iodide  (p.  278). 

Phosphorus  trichloride  PC13  is  made  by  passing  chlorine  gas  over 
melted  phosphorus  in  a  flask  until  the  proper  gain  in  weight  has 
occurred;  It  is  a  liquid,  boiling  at  76°.  When  excess  of  chlorine  is 
employed,  phosphorus  pentachloride  PCU,  which  is  a  white  solid 
body,  is  formed.  When  moist  air  is  blown  over  any  of  these  sub- 
tances,  the  water  is  condensed  to  a  fog  by  the  hydrogen  halide.  In 


PHOSPHORUS  555 

the  case  of  the  interaction  of  phosphorus  pentachloride  and  water, 
phosphoric  acid  is  formed: 

PC15  +  4H20  -> H3P04  +  5HC1. 

With  a  limited  supply  of  water  the  hydrolysis  is  not  so  complete, 
and  phosphoryl  chloride*  (phosphorus  oxy chloride),  a  liquid  boiling 
at  107°,  is  produced:  PC15  +  HOH  -*  POC13  +  2HC1. 

This  interaction  of  phosphorus  pentachloride  and  water  is  a  per- 
fectly general  one,  and  takes  place  with  most  compounds  containing 
hydroxyl.  Thus,  when  alcohol  (which  differs  from  water  in  having 
ethyl  C2H5,  instead  of  hydrogen,  combined  with  hydroxyl)  is  poured 
upon  it,  ethyl  chloride,  phosphoryl  chloride,  and  hydrogen  chloride 
are  formed: 

C2H5OH  +  PC15  ->  C2H6C1  +  POC13  +  HC1. 

The  same  action  takes  place  with  all  carbon  compounds  containing 
hydroxyl,  and  is  used  as  a  means  of  showing  the  presence  of  this 
group  in  their  structure.  The  reaction  is  shown  by  inorganic  com- 
pounds also.  Thus,  anhydrous  sulphuric  acid  gives  sulphuryl 
chloride,  which  may  be  separated  from  the  phosphoryl  chloride  by 
fractional  distillation  (see  Petroleum)  : 

S02(OH)2  +  2PC15  ->  S02C12  +  2POC13  +  2HC1. 

Phosphorus  pentachloride,  when  heated,  reaches  a  vapor  tension 
of  760  mm.  at  163°,  and  while  still  solid.  It  therefore  passes  freely 
into  vapor  (boils,  so  to  speak)  at  this  temperature,  and  condenses  di- 
rectly to  the  solid  form.  This  sort  of  distillation  is  called  sublimation. 
At  a  pressure  above  that  of  the  atmosphere  it  melts  at  166°.  This  is 
simply  a  case  in  which  the  vapor  tension  of  the  solid,  increasing  with 
rise  in  temperature,  happens  to  pass  the  arbitrary  value  of  the  oppos- 
ing pressure  (one  atmosphere)  peculiar  to  experiments  carried  on  in 
open  vessels,  before  the  melting-point  is  reached.  The  same  phe- 
nomenon is  shown  by  sulphur  trioxide  (p.  430). 

Phosphorus  pentachloride  (cf.  p.  260)  and  pentabromide,  when 
vaporized,  are  partially  dissociated: 

PBr6  <=±  PBr3  +  Br2. 

Since  the  first  two  members  of  this  equilibrium  are  colorless,  while  the 
bromine  is  brown,  this  action  may  be  used  to  illustrate  the  effect  upon 

*  This  substance  is  a  mixed  anhydride  (p.  449)  of  phosphoric  acid  and  hydro 
gen  chloride. 


556  INORGANIC  CHEMISTRY 

a  system,  of  increasing  the  concentration  of  one  of  the  interacting 
substances  (p.  291).  Two  tubes  of  equal  volume  and  containing 
equal  amounts  of  the  pentabromide  are  prepared.  A  small  amount  of 
the  tribromide  is  added  to  the  second,  and  both  are  sealed  up.  When 
the  tubes  are  now  heated  to  the  same  temperature,  the  contents  of  the 
second  will  be  less  strongly  colored  by  bromine  in  consequence  of  the 
greater  activity  in  it  of  the  reversing  action.  At  163°  and  760  mm., 
about  4  per  cent  of  the  molecules  of  the  vapor  of  the  pentachloride 
are  dissociated  into  the  trichloride  and  chlorine. 

Oxides  of  Phosphorus.  —  The  oxides  of  phosphorus  are  the  so- 
called  trioxide  PA,  the  pentoxide  P2C>5,  and  a  tetroxide  P2O4. 

The  pentoxide  is  a  white  powder  formed  when  phosphorus  is 
burned  with  a  free  supply  of  oxygen.  It  unites  with  water  with 
great  violence  to  form  metaphosphoric  acid  (see  below),  and  hence  is 
known  as  phosphoric  anhydride:  P205  +  H2O  — >  2HPO3.  In  the 
laboratory  this  action  is  frequently  utilized  for  drying  gases  (p.  123) 
and  for  removing  water  from  combination  (p.  527).  The  vapor 
density  of  the  pentoxide  indicates  that  its  formula  is  P^io,  use  of 
which,  however,  would  only  complicate  our  equations. 

The  trioxide  P4Oe  is  obtained  by  burning  phosphorus  in  a  tube 
with  a  restricted  supply  of  air.  It  is  a  white  solid,  melting  at  22.5° 
and  boiling  at  173°.  On  account  of  the  ease  with  which  it  may  be 
volatilized,  it  can  be  separated  by  distillation  from  any  pentoxide 
formed  at  the  same  time.  The  operation  must  be  carried  out  in  an 
apparatus  from  which  the  air  is  excluded,  as  the  trioxide  unites 
spontaneously  with  oxygen.  The  vapor  is  phosphorescent  (p.  550). 
The  vapor  density  of  the  substance  shows  that  its  formula  is  P40e. 
This  formula  is  preferred  to  the  simpler  one  because,  although  the 
oxide  is  the  anhydride  of  phosphorous  acid,  it  nevertheless  unites 
exceedingly  slowly  with  cold  water  to  form  this  substance.  It  inter- 
acts vigorously  with  hot  water,  but  phosphine,  red  phosphorus,  hypo- 
phosphoric  acid,  and  phosphoric  acid  are  amongst  the  products,  and 
very  little  phosphorous  acid  escapes  decomposition.  When  this  oxide 
is  heated  to  440°  it  decomposes,  giving  the  tetroxide  P2O4  and  red 
phosphorus. 

Acids  of  Phosphorus.  —  There  are  six  different  acids  of 
phosphorus  in  which  four  distinct  stages  of  oxidation  are  shown. 
The  highest  stage  is  represented  by  three  phosphoric  acids,  where 
the  degree  of  hydration  of  the  anhydride  varies.  The  others  show 


PHOSPHORUS  557 

three  different  and  lower  states  of  oxidation  (study  their  positive  and 
negative  valences,  p.  494)  : 

Orthophosphoric  acid  H3PO4    (  =  3H2O,P2O6) 

Pyrophosphoric  acid  H4P2O7  (  =  2H2O,P2O6) 

Metaphosphoric  acid  HPO3     (=  H2O,P2O6) 

Hypophosphoric  acid  H2PO3    (  =  2H2O,P2O4) 

Phosphorous  acid  H3PO3    (=  3H2O,P2O3) 

Hypophosphorous  acid  H3PO2   (=  3H2O,P2O) 

The  Phosphoric  Acids.  —  The  relation  between  the  three  differ- 
ent phosphoric  acids  may  be  seen  by  considering  them  as  being 
formed  from  phosphorus  pentoxide  and  water.  It  will  be  remembered 
that  in  the  majority  of  cases  already  considered,  this  sort  of  action 
takes  place  for  the  most  part  in  but  one  way.  Thus,  nitric  acid  is 
known  in  but  one  form,  which  is  produced  by  the  union  of  one  molecule 
each  of  nitrogen  pentoxide  and  water:  N2O5  +  H2O  —  >  2HN03. 
Similarly  the  chief  sulphuric  acid  is  the  one  formed  from  one  molecule 
of  sulphur  trioxide  and  one  molecule  of  water:  S03  +  H2O  —  >  H2SO4, 
although  here  we  have  both  the  hydrate  H2SO4,H20,  which  might 
be  written  H4SO5  and  disulphuric  acid  H2S207.  Referred  to  the 
anhydride  these  three  acids  are  H2O,S03,  2H2O,S03,  and  H2O,2S03. 
Periodic  acid  (p.  487)  has  a  set  of  even  more  complexly  related  acids 
or  salts. 

Now,  when  phosphoric  anhydride  acts  upon  water  we  obtain  a 
solution  which,  on  immediate  evaporation,  leaves  a  glassy  solid, 
HPO3,  known  as  metaphosphoric  acid.  This  is  H2O,P205.  When, 
however,  the  solution  is  allowed  to  stand  for  some  days,  or  is  boiled 
with  a  little  dilute  nitric  acid,  the  hydrogen-ion  of  which  acts  catalyt- 
ically,  the  residue  from  evaporation  is  H3PO4,  Orthophosphoric  acid 
3H2O,P2O6: 

P205  +  3H20  -»  2H3P04    or    HPO3  +  H2O  -»  HsPO^ 

Conversely,  when  Orthophosphoric  acid  is  kept  at  about  255°  for  a 
time,  it  slowly  loses  water,  and  H4P2O7,  pyrophosphoric  acid,  is  ob- 
tained : 

2H3P04  -»  H^OT  +  H20  T  • 


This  acid  is  2H2O,P2C>5.  Further  desiccation  leaves  metaphosphoric 
acid,  which  cannot  be  further  resolved  into  phosphorus  pentoxide  and 
water.  When  dissolved  in  water,  pyrophosphoric  acid  slowly  resumes 
the  water  which  it  has  lost  and  gives  the  prtho-acid  again. 


558  INORGANIC  CHEMISTRY 

The  relations  of  all  these  substances  are  more  clearly  seen  in  the 
graphic  formulae: 

=  0 

-0-H 

-0-H         ,J  -0-H  Pi  -0-H  ^f=0 

-0-H 

-0-H 
-0-H 
-0-H 


-O-H 

0          ->  <-       J     0 

-0-H  f  =0  p     =0 

-0-H  P\  -0-H  M  =0 

=  0 


=o 

Ortho-  Pyro-  Meta-  Oxide 

The  addition  or  removal  of  water  leaves  the  valence,  and  therefore 
the  degree  of  oxidation,  of  the  phosphorus  unchanged. 

Pyrosulphuric  acid  and  its  salts  when  dissolved  in  water  give  sul- 
phuric acid  and  acid  sulphates,  respectively.  That  is  to  say,  the  ion 
820?  is  not  capable  of  existence.  But  the  very  slow  rate  at  which 
the  less  hydrated  phosphoric  _acids  change  into  the  more  hydrated 
ones  shows  that  ions  like  PO4— ,  POs~,  and  P2O7=~may  be  compara- 
tively stable.  The  behavior  of  solutions  of  the  salts  shows  this  even 
more  clearly. 

Qrthophosphoric  Acid  HSPO^.  —  As  we  have  seen  (p.  548), 
ordinary  calcium  phosphate  is  the  source  of  the  impure,  commercial 
acid.  Pure  orthophosphoric  acid  may  be  made  by  boiling  red 
phosphorus  with  slightly  diluted  nitric  acid  and  evaporating  off  the 
water  and  excess  of  nitric  acid.  The  product  is  a  white,  crystalline, 
deliquescent  hemihydrate,  2H3PO4,H20  (m.-p.  29.35°).  Anhydrous 
orthophosphoric  acid  melts  at  42.3°. 

This  acid  is  much  weaker  than  sulphuric  acid,  and  is  dissociated 
chiefly  into  the  ions  H+  and  H2P04~.  The  dihydrophosphate-ion  is 
broken  up  to  some  extent  into  H+  and  HP04=,  as  we  learn  from  the 
fact  that  the  solution  of  the  sodium  salt  NaH2P04  is  acid.  The  ion 
HPO4=  is  hardly  dissociated  at  all,  for  a  solution  of  the  salt  Na2HP04 
is  not  acid  in  reaction. 

Salts  of  Orthophosphoric  Acid.  —  As  a  tribasic  acid,  it 
forms  salts  of  three  kinds,  such  as  NaH2PO4,  Na2HPO4,  and  Na3P04. 
These  are  known  respectively  as  primary,  secondary,  and  tertiary 
sodium  ortbophosphate.  The  primary  sodium  phosphate  is  faintly 


PHOSPHORUS  559 

acid  in  reaction.  The  secondary  one  is  slightly  alkaline,  because  of 
hydrolysis  arising  from  the  tendency  of  the  hydrogen-ion  of  the 
water  to  combine  with  the  HP04~  to  form  H2PO4~,  which  is  much  more 
feebly  acid  than  is  phosphoric  acid  H3P04.  The  simplified  equation 
(p.  419)  shows  the  reason  for  the  alkalinity  of  the  solution:  HPO4=  + 
H+  +  OH~  <=>  H2PO4~  +  OH~,  for  hydroxyl-ion  is  present.  The 
tertiary  phosphate  is  stable  only  in  solid  form,  and  can  be  made  by 
evaporating  to  dryness  a  mixture  of  the  secondary  phosphate  and 
sodium  hydroxide: 

Na*HP04  +  NaOH  <=»  Na3P04  +  H20|. 

When  the  product  is  dissolved  in  water,  the  action  is  reversed  (cf. 
p.  418).  Mixed  phosphates  are  also  known,  particularly  sodium- 
ammonium  phosphate  (microcosmic  salt)  NaNH4HPO4,  and  the 
insoluble  magnesium-ammonium  phosphate,  MgNH4PO4.  . 

Primary  calcium  phosphate,  known  in  commerce  as  superphos- 
phate, is  used  as  a  fertilizer.  Since  plants  can  take  up  soluble 
substances  only,  the  insoluble,  natural  calcium  phosphate  is  of 
relatively  little  service  to  plants.  It  is  therefore  converted  into  the 
superphosphate,  which  is  soluble,  by  treatment  with  dilute  sulphuric 
acid: 

Ca3(P04)2  +  2H2SO4  *±  2CaS04  +  CaH4(PO4)2. 

The  tertiary  phosphates  are  unchanged  by  heating.  The  primary 
and  secondary  phosphates,  however,  retaining,  as  they  do,  some  of 
the  original  hydrogen  of  the  phosphoric  acid,  are  capable  of  losing 
water  like  phosphoric  acid  itself,  when  heated.  The  actions  are 
slowly  reversed  when  the  products  are  dissolved  in  water: 

NaH2P04  <=»  NaP03  +  H20, 
2Na*HPO4  <=±  Na4P207  +  H2O. 

It  will  be  seen  that  the  meta-  and  pyrophosphates  of  sodium  are 
formed  by  these  actions;  and  this  is  indeed  the  simplest  way  of  form- 
ing these  substances,  since  the  acids  themselves  are  not  permanent  in 
solution,  and  are  too  feeble  to  lend  themselves  to  exact  neutralization. 
Ammonium  salts  of  phosphoric  acid  lose  ammonia,  as  well  as  water, 
when  heated  (cf.  p.  520).  Thus,  microcosmic  salt  gives  first  the 
primary  sodium  phosphate: 

NaNH4HP04  -> NH3T  +  NaH2P04  ->  NaP03  +  H20t, 
and  this  in  turn  loses  water  to  give  the  metaphosphate. 


560  INORGANIC   CHEMISTRY 

Pyrophosphoric  Acid  H4P2(>7.  —  This  acid,  obtained  by  heat- 
ing orthophosphoric  acid,  may  be  prepared  in  pure  form  by  making 
the  sparingly  soluble  lead  salt  from  sodium  pyrophosphate,  and 
precipitating  the  lead-ion  as  lead  sulphate  by  addition  of  sulphuric 
acid.  In  solution  it  gradually  reunites  with  water.  Although 
tetrabasic,  having  four  hydrogen  atoms  which  may  be  displaced  by 
metals,  only  two  kinds  of  salts  are  known.  These  are  the  normal 
salts,  such  as  Na4P2O7,  and  those  in  which  one-half  of  the  hydrogen 
has  been  displaced  by  a  metal,  such  as  Na2H2P207. 

Metaphosphoric  Acid  HPOS.  —  This  is  the  "  glacial  phosphoric 
acid"  of  commerce,  and  is  usually  sold  in  the  form  of  transparent 
sticks.  It  is  obtained  by  heating  orthophosphoric  acid,  or  by  direct 
union  of  phosphorus  pentoxide  with  a  small  amount  of  cold  water. 
It  passes,  into  vapor  at  a  high  temperature,  and  its  vapor  density 
corresponds  to  the  formula  (HP03)2.  The  existence  of  certain  com- 
plex salts  confirms  our  belief  in  the  existence  of  association  (p.  282). 

Sodium  metaphosphate  NaPO3,  in  the  form  of  a  small  globule 
obtained  by  heating  microcosmic  salt  on  a  platinum  wire,  is  used  in 
analysis.  When  minute  traces  of  oxides  of  certain  metals  are  placed 
upon  such  a  globule,  known  as  a  bead,  and  heated  in  the  Bunsen 
flame,  the  mass  is  colored  in  various  tints  according  to  the  oxide 
used  (bead  test).  This  action  may  be  understood  when  we  consider 
that  sodium  metaphosphate  takes  up  water  to  form  primary  sodium 
orthophosphate:  NaP03  +  H2O  — »  NaH2P04.  In  the  same  way,  but 
at  higher  temperatures,  it  is  able  to  take  up  oxides  of  elements  other 
than  hydrogen,  giving  mixed  orthophosphates.  Thus,  with  oxide  of 
cobalt,  a  part  of  the  metaphosphate  unites  according  to  the  equation : 

NaP03  +  CoO  -»  NaCoPO4, 
and  the  product  gives  a  blue  color  to  the  bead. 

Distinguishing  Tests.  —  When  a  solution  of  nitrate  of  silver  is 
added  to  a  solution  of  orthophosphoric  acid  or  any  soluble  orthophos- 
phate, a  yellow  precipitate  of  silver  orthophosphate  Ag3PO4  is  pro- 
duced. This  is  a  test  for  orthophosphate-ion.  With  pyrophosphoric 
acid  or  any  pyrophosphate  the  product  is  white  Ag4P2C>7.  With 
metaphosphoric  acid  a  white  precipitate,  AgP03,  is  obtained  also. 
Metaphosphoric  acid  coagulates  a  clear  solution  (really  a  colloidal 
suspension)  of  albumen  (say,  white  of  egg),  while  ortho-  or  pyro- 
phosphoric acid  has  no  visible  effect  upon  it. 


PHOSPHORUS  561 

A  test  for  orthophosphoric  acid,  or  rather  the  ion  PO^,  consists 
in  adding  a  drop  of  the  solution  containing  this  ion  to  a  solution  of 
ammonium  molybdate  (q.v.)  in  dilute  nitric  acid.  A  copious  yellow 
precipitate  of  an  ammonium  phosphomolybdate  (NH^sPO^llMo- 
O3,6H2O  appears  on  warming.  In  presence  of  excess  of  ammonia,  the 
formation  of  the  white  insoluble  ammonium-magnesium  phosphate 
(p.  559)  serves  as  a  test  also.  Arsenic  acid  (q.v.)  gives  precipitates 
of  appearance  and  composition  similar  to  these  two. 

Phosphorous  Acid  HsPOz.  —  With  cold  water  phosphorus  tri- 
oxide  P4Oe  yields  phosphorous  acid  very  slowly.  With  hot  water  the 
action  is  exceedingly  violent  and  complex  (p.  556).  This  acid  may 
be  obtained  also  by  the  action  of  water  upon  phosphorus  trichloride, 
tribromide,  or  tri-iodide  and  evaporation  of  the  solution : 

PC13  +  3H2O  ->  P(OH)3  +  3HC1. 

Some  of  this  acid,  along  with  phosphoric  acid  and  hypophosphoric 
acid,  is  formed  when  moist  phosphorus  oxidizes  in  the  air. 

In  spite  of  the  presence  of  three  hydrogen  atoms,  this  acid  is 
dibasic,  and  two  only  are  replaceable  by  metals.  To  express  this  fact, 
the  first  of  the  following  formulae  is  preferred: 


since  the  symmetrical  formula  would  indicate  no  difference  between 
the  three  hydrogen  atoms.  H  united  directly  to  P,  as  here  and  in 
PH3,  is  practically  not  acidic.  Phosphorous  acid  is  a  powerful 
reducing  agent,  precipitating  silver,  for  example,  in  the  metallic 
form  from  solutions  of  its  salts.  When  heated,  it  decomposes,  giving 
the  most  stable  acid  of  phosphorus  (cf.  pp.  447,  474,  483),  namely, 
metaphosphoric  acid,  and  phosphine: 

4H3P03  -»  3HP03  +  3H2O  +  PH3|. 

Hypophosphorous  Acid.  —  The  potassium  salt  of  this  acid  is 
obtained,  as  we  have  seen,  when  phosphorus  is  heated  with  potassium 
hydroxide  solution  (p.  553) .  It  may  be  prepared  in  the  free  form  by 
substituting  barium  hydroxide  for  potassium  hydroxide: 

3Ba(OH)2  +  8P  +  6H2O  -»  3Ba(H2PO2)2  +  2PHat, 


562  INORGANIC  CHEMISTRY 

By  careful  addition  of  dilute  sulphuric  acid  to  the  resulting  liquid, 
barium  sulphate  is  precipitated.  On  evaporation  of  the  water,  the 
white  crystalline  acid  H3PO2  is  obtained.  This  acid  is  monobasic; 
two  of  its  hydrogen  atoms  cannot  be  displaced  by  metals.  To  ex- 

r-H 

press  this  fact  the  graphic  formula  0  =  P  •<  —  H   is  used.    This  sub- 

(-OH 

stance  is  also  a  powerful  reducing  agent,  tending,  by  the  acquisition  of 
oxygen,  to  pass  into  phosphoric  acid. 

The  acid  and  its  salts,  when  heated,  give  off  water  and  phosphine, 
and  leave  phosphoric  acid  and  phosphates  (cf.  p.  561),  respectively. 
Thus,  sodium  hypophosphite  gives  phosphine  and,  finally,  sodium 
pyrophosphate : 

4NaH2P02  i  2PH3T  +  2Na2HPO4  ->  Na4P207  +  H2O| . 

Hypophosphoric  Acid  /f2PO3.  —  When  white  phosphorus  is 
heated  under  nitric  acid,  alone,  orthophosphoric  acid  is  formed. 
When,  however,  cupric  nitrate  (or  silver  nitrate;  other  nitrates  are 
without  effect)  is  added,  copper  is  precipitated  (as  Cu3P2  or  Gu). 
By  neutralizing  half  of  the  resulting  solution  with  sodium  carbonate, 
and  adding  the  other  half,  a  copious  precipitate  of  a  hydrate  of  sodium 
hypophosphate  NaHPO3,2H2O  appears.  From  this  salt,  other  salts 
and  the  free  acid  can  be  prepared. 

Structural  Formulae  of  Salts  of  Hydrogen.  —  As  a  rule,  the 
formulae  of  acids  have  thus  far  been  written  with  the  ionizable  hydro- 
gen in  front:  HC1,  H2SO4,  HCO2CH3.  This  is  only  one  illustration  of 
the  method  by  which  chemists  have  constantly  sought  to  utilize  for- 
mulae for  the  purpose  of  expressing,  not  merely  the  composition  of  a 
substance,  but  some  of  its  properties  as  well.  By  another  typo- 
graphical device  we  have  attempted  to  indicate  the  behavior  of  dilute 
solutions  by  putting  the  radicals  in  brackets:  Cu(NO3)2,  Ba(OH)2. 
These  are  called  reaction  formulae,  and  their  object  is  to  exhibit  the 
modes  of  action  of  the  substance.  Now  the  modes  of  action  of  a 
single  substance  are  often  rather  various,  and  one  and  the  same 
structural  formula  cannot  represent  all  of  these  at  once.  We  have 
observed  this,  particularly,  with  the  oxygen  acids.  Thus,  H2.S04 
expresses  the  mode  of  activity  in  dilute  solution  and  often  when 
no  solvent  is  present,  as  in  the  action  on  chlorides  (p.  206)  and 
nitrates  (p.  526).  But  when  all  the  hydrogen  of  an  acid  is  not  ioniz- 


PHOSPHORUS  563 

able,  we  regard  that  which  is  so  as  part  of  an  hydroxyl  group  in  the 
parent  molecules,  and  the  rest  as  being  attached  to  the  characteristic 
non-metal  of  the  acid,  as,  for  example,  in  phosphorous  acid  (p.  561; 
cf.  pp.  442,  540).  Thus,  we  should  write  phosphorous  acid  HPO(OH)2, 
instead  of  H2PO3H,  to  chronicle  this  fact.  So  also  the  formula 
POH2(OH)  is  used  for  hypophosphorous  acid.  Molecular  actions, 
such  as  those  of  sulphuric  acid  S02(OH)2  (p.  272),  are  well  shown  by 
these  formulae: 

S02(OH)2  +  2HBr  ->  S02  +  2H20  +  Br2. 

It  must  be  noted  particularly  that  this  sort  of  formula,  when  the  sub- 
stance for  which  it  stands  is  an  acid,  represents  only  some  features 
in  the  behavior  of  the  anhydrous  substance  and  of  the  molecules,  and 
not  the  ionic  action  in  solution.  A  formula  like  Ba(OH)2,  where  the 
material  is  a  base,  on  the  other  hand,  represents  both  the  ionic  and  the 
molecular  behavior.  The  graphic  formula  is  more  general  (cf.  p.  540). 
It  shows  all  these  relations,  and  often  still  others,  but  none  of  them  so 
specifically. 

Sulphides  of  Phosphorus.  —  White  phosphorus,  when  heated 
with  sulphur  unites  with  explosive  violence.  By  using  red  phos- 
phorus the  action  can  be  controlled.  By  employing  the  proper 
proportions,  the  pentasulphide  P2$5  is  secured.  It  is  purified  by 
distillation  from  a  retort  in  which  a  current  of  carbon  dioxide  is 
maintained  (see  below) .  The  distillate  solidifies  to  a  yellow,  crystal- 
line solid  (m.-p.  290°,  b.-p.  515°).  Materials  undergoing  chemical 
change,  which  are  to  be  kept  at  a  constant,  high  temperature,  are 
often  placed  in  tubes  suspended  in  the  vapor  of  the  pentasulphide. 
When  a  lower  temperature  is  required,  boiling  sulphur  (445°)  is  used. 

Distillation  in  a  stream  of  some  inactive  gas  is  a  common  means 
of  distilling  under  reduced  pressure  (cf.  p.  316).  The  dilution  of  the 
vapor  lowers  its  partial  pressure,  just  as  would  evacuation.  This  plan 
has  the  advantage,  however,  of  sweeping  the  vapor  away  from  the 
heated  region  into  the  condenser,  and  so  diminishing  the  amount  of 
decomposition.  In  dealing  with  compounds  of  carbon,  a  current  of 
steam  is  often  used  for  the  above  purposes.  It  enables  us  also  to  sepa- 
rate a  slightly  volatile  substance  from  one  which  is  almost  involatile. 

Phosphorus  pentasulphide  is  hydrolyzed  by  cold  water,  and  acts 
upon  other  substances  containing  hydroxyl  when  heated  with  them,, 
the  actions  being  similar  to  those  of  the  pentachloride  (p.  555) : 

P2S5  +  8H20  -»  2H3P04  +  5H2S. 


564  INORGANIC  CHEMISTRY 


Other  sulphides,  P4Sa  (used  in  making  matches,  p.  552),  and 
P4S7,  may  be  prepared  by  using  the  constituents  in  the  proportions 
represented  by  these  formulae. 

Comparison  of  Phosphorus  with  Nitrogen  and  with  Sul- 
phur. —  Although  phosphorus  and  nitrogen  are  regarded  as  belong- 
ing to  one  family,  the  differences  between  them  are  more  conspicuous 
than  the  resemblances.  The  latter  are  confined  almost  wholly  to 
matters  concerned  with  valence.  The  differences  are  seen  in  the 
facts  that  nitrogen  is  a  gas,  while  phosphorus  is  a  solid  occurring 
in  two  varieties,  and  that  the  former  is  inactive  and  the  latter  active. 
The  contrasts  between  phosphine  and  ammonia  (pp.  553-554)  and 
between  the  halides  of  the  two  elements  (p.  554)  have  been  noted 
already.  The  pentoxide  of  nitrogen  decomposes  spontaneously;  that 
of  phosphorus  is  one  of  the  most  stable  of  compounds.  Nitric  acid 
is  very  active,  both  as  an  acid,  and  as  an  oxidizing  agent;  the  phos- 
phoric acids  are  quite  the  reverse. 

On  the  other  hand,  the  resemblance  of  phosphorus  to  sulphur  is 
marked.  Both  are  solids,  existing  in  several  forms.  Both  yield 
stable  compounds  with  oxygen  and  chlorine.  The  hydrogen  com- 
pounds interact  with  salts  to  give  phosphides  of  metals  and  sulphides 
of  metals,  respectively.  Against  these  must  be  set  the  facts,  that 
hydrogen  sulphide  does  not  unite  with  the  hydrogen  halides  at  all, 
while  phosphine  gives  the  phosphonium  halides,  and  that  phosphoric 
acid  is  hard  to  reduce  while  sulphuric  acid  is  reduced  with  compara- 
tive ease. 

Exercises.  —  1.  Explain  the  effect  of  sulphuric  acid  in  setting 
fire  to  the  earliest  matches  (p.  552). 

2.  Make   a   brief    definition    of   a    substance   which    sublimes 
(p.  555). 

3.  Why  would  a  mixture  of  potassium  dichromate  and  hydro- 
chloric acid  (p.  418)  be  less  suitable  than  nitric  acid  for  making 
phosphoric  acid  from  red  phosphorus? 

4.  Why  is  not  the  tertiary  phosphate  of  sodium  (p.  559)  decom- 
posed by  heating?    What  tertiary  phosphates  would  be  decomposed 
by  this  means? 

5.  Formulate   the   hydrolyses   of   the   secondary   and   tertiary 
sodium  orthophosphates  as  was  done  for  sodium  sulphide  (p.  418). 

6.  How  should  you  prepare  Ca2P2O7  and  Ca(POs)2? 

7.  What  product  should  you  confidently  expect  to  find  after 


PHOSPHORUS  565 

heating,  (a)  sodium  phosphite  Na2HPO3,  (b)  potassium  hypophosphite 
(p.  561)?     Make  the  equations. 

8.  Compare  the  elements  chlorine  and  phosphorus  after  the  man- 
ner of  the  comparisons  on  p.  564. 

9.  What  are  the  valences  of  the  non-metals  in:  H2S207,  H2Cr207, 
KMnO4,   KH2PO2,  H3NO4,  NaH2PO3,  Na2P03?     Name  these  sub- 
stances. 

10.  Is  it  oxidation  or  reduction,  or  neither,  when  we  make,  (a) 
N204  from  HNO3,  (6)  SO2  from  H2SO3,  (c)  HPO3  from  H3PO3,  (d) 
H2S2O7  from  H2S04,  (e)  Na2SO4  from  NaHSO3? 


CHAPTER    XXVIII 
CARBON  AND   THE   OXIDES   OF   CARBON 

The  Chemical  Relations  of  the  Element.  —  The  elements  of 
the  carbon  family  are  carbon,  silicon,  germanium,  tin,  and  lead 
(see  Periodic  system).  Of  these  the  first  two  are  entirely  non- 
metallic,  while  the  others  are  metallic  elements  showing  more  or  less 
strong  resemblances  to  the  non-metals.  All  these  elements  are 
quadrivalent  as  regards  the  maximum  valence  which  they  exhibit. 
With  the  exception  of  silicon,  however,  they  all  form  many  com- 
pounds in  which  they  are  bivalent. 

The  chemistry  of  the  compounds,  of  carbon  is  an  exceedingly  ex- 
tensive and  complex  subject.  It  is  commonly  known  as  organic 
chemistry,  on  account  of  the  fact  that  the  majority  of  the  substances 
composing,  and  produced  by,  living  organisms  are  compounds  of 
carbon,  and  that  it  was  at  first  supposed  that  their  artificial  produc- 
tion, e.g.,  without  the  intervention  of  life,  was  impossible.  But  many 
natural  organic  products  have  now  been  made  from  simpler  ones  or 
from  the  elements,  a  process  called  synthesis,  and  the  preparation  of 
the  others  is  delayed  only  in  consequence  of  difficulties  caused  by 
their  instability  and  complexity.  On  the  other  hand,  thousands  of 
carbon  compounds,  unknown  to  animal  or  vegetable  life,  including 
many  valuable  drugs  and  dyes,  have  now  been  added  to  the  cata- 
logue of  chemical  compounds.  More  than  200,000  different  com- 
pounds containing  carbon  are  known,  and  thousands  are  added 
every  year. 

The  elements  entering  into  carbon  compounds  are  chiefly  hydrogen 
and  oxygen.  After  these  come  nitrogen,  the  halogens,  and  sulphur. 

CARBON 

Occurrence.  —  Large  quantities  of  carbon  are  found  in  the  free 
condition  in  nature.  The  diamond  is  the  purest  natural  carbon,  and 
at  the  same  time  the  least  plentiful.  Graphite,  or  plumbago,  which  is 
the  next  purest,  is  found  in  limited  amounts,  and  is  a  valuable  mineral. 
Coal  occurs  in  numerous  forms,  but  much  of  it  contains  no  free  carbon. 
Small  quantities  of  the  free  element  have  been  found  in  meteorites. 

566 


CARBON  AND  THE  OXIDES  OF  CARBON 


567 


In  combination,  carbon  is  found  in  marsh-gas,  or  methane  CH4, 
which  is  the  chief  component  of  natural  gas.  The  mineral  oils  consist 
almost  entirely  of  mixtures  of  various  compounds  of  carbon  and 
hydrogen.  Whole  geological  formations  are  composed  of  carbonates 
of  common  metals,  particularly  calcium  carbonate  or  limestone,  and 
a  double  carbonate  of  calcium  and  magnesium,  known  as  dolomite. 

Allotropic  Forms  of  Carbon.  —  The  allotropic  (p.  315)  forms 
of  carbon  differ  very  strikingly  in  their  physical  properties.  The 
diamond  (density  3.5)  is  transparent,  crystalline,  and  very  hard. 
Graphite  (density  2.3)  is  black,  lustrous,  and  very  soft.  Amorphous 
carbon  is  very  variable.  Thus  lampblack  (see  p.  596)  is  a  fine 
powder  of  nearly  pure  carbon,  charcoal  (see  p.  610)  shows  the 
structure  of  the  wood.  These  amorphous  forms  can  best  be  dis- 
cussed after  the  materials  from  which  they  are  formed  have  been 
considered. 

That  all  the  forms  are  composed  of  the  same  element  is  shown 
by  the  fact  that  they  all  burn  in  oxygen  to  give  carbon  dioxide. 
Then,  too,  when  heated  strongly  in  absence  of  air,  diamond  and 
the  amorphous  forms  all  turn  into  graphite.  They  contain  differ- 
ent amounts  of  chemical  energy,  however.  Thus,  when  1  g.  of 
each  is  burned,  diamond  gives  7870  cal.,  graphite  7835  and  sugar 
charcoal  (p.  438)  8040.  The  tendency  of  most  carbon  compounds, 
when  heated,  to  char,  giving  free  carbon,  is  used  as  a  test. 

The  Diamond.  —  Diamonds,  which  are  found  chiefly  in  Brazil 
and  South  Africa,  are  scattered  sparsely  through 
metamorphic  and  volcanic  rocks  which  seem  to 
have  undergone  secondary  changes.  They  are 
covered  with  a  crust  which  entirely  obscures  their 
luster,  and  possess  natural  crystalline  forms  be- 
longing to  the  regular  system.  A  form  related  to 
the  octahedron  (p.  172)  is  frequently  observed. 
It  should  be  noted  that  this  natural  form  bears  no 
relation  whatever  to  the  pseudo-crystalline  shape 
which  is  conferred  upon  the  stone  by  the  diamond- 
cutter.  Thus,  a  " brilliant"  possesses  one  rather 

large,  flat  face,  which  forms  the  base  of  a  many    

sided  pyramid  (Fig.  126,  showing  two  views).    This  FlQ  126 

form  is  given  to  the  stone,  in  order  that  the  max- 
imum reflection  of  light  from  its  interior  may  be  produced.     The 


568  INORGANIC  CHEMISTRY 

diamond  is  harder  (see  Appendix  II)  than  any  other  variety  of 
matter,  with  the  exception,  perhaps,  of  one  carbide  of  boron,  while 
only  one  or  two  other  materials,  like  carborundum,  approach  it. 
Hence,  it  can  be  scratched  or  polished  only  by  rubbing  with 
diamond  powder.  It  is  the  densest  form  of  carbon  (density  3.5). 
The  colorless  stones,  and  occasional  specimens  with  special  tints 
(like  the  blue,  Hope  diamond)  are  the  most  valuable.  The  black 
("carbonado")  and  discolored  specimens  are  used  for  grinding  and 
glass  cutting.  Mounted  round  the  edge  of  a  tube,  they  are  used 
for  drilling  rock,  so  that  a  cylindrical  specimen  of  the  whole  of  the 
strata  can  be  secured  for  examination.  All  forms  of  carbon  are 
insoluble  in  all  liquids  at  room  temperature.  Molten  iron  (q.v.) 
dissolves  five  or  six  per  cent,  part  of  which  goes  into  combination. 
The  diamond  is  a  nonconductor  of  electricity.  Diamonds  are  sold 
by  the  new  international  carat,  200  mgms.  (old  carat,  4  grains  = 
205  mgms.),  and  the  value  increases  with  the  size.  Thus,  a  first 
quality,  cut  stone  of  1  carat  is  worth  about  $270,  one  of  2  carats 
about  $340  per  carat.  The  largest  diamond  known,  the  Cullinan 
(1905),  weighed  3032  (old)  carats  (621  g.  or  1.37  Ibs).  It  was  pre- 
sented by  the  Transvaal  government  to  King  Edward  VII,  and  was 
cut  into  stones  of  516.5  and  309  carats  and  many  smalley  ones. 
Other  large  stones  are  the  Nizam  (277  carats),  the  Jubilee  (239  carats), 
and  the  Kohinoor  (106  carats). 

The  diamond,  although  its  origin  in  nature  is  still  a  matter  of  uncer- 
tainty, has  been  made  artificially.  Moissan  (1887)  dissolved  carbon 
in  molten  iron  and,  after  chilling  the  mass  so  as  to  produce  a  solid 
crust,  which  by  its  shrinkage  severely  compressed  the  interior, 
allowed  the  whole  to  cool  very  slowly.  Portions  of  the  interior 
of  the  ingot  were  treated  with  acid  to  dissolve  the  iron,  and  amongst 
the  insoluble  particles  were  recognized  a  few  microscopic  fragments 
(none  larger  than  0.5  mm.)  which  exhibited  the  form  and  hardness 
of  the  diamond.  The  greater  part  of  the  carbon,  however,  appeared, 
as  usual,  as  graphite. 

Graphite.  —  Graphite  (Gk.  >7>a<£a>,  I  write)  or  plumbago  is 
found  in  Cumberland,  Siberia,  Ceylon  (1913,  28,500  short  tons), 
Canada,  and  Austria  (1912,  50,000  short  tons).  It  is  composed  of 
glittering,  slippery  scales.  Good  crystals  are  seldom  found  (hexagonal 
system).  The  mineral  is  extremely  soft,  in  utter  contrast  to  the 
diamond,  and  has  a  lower  density  (2.3).  It  also  conducts  electricity. 
It  is  now  made  artificially  by  an  electro-thermal  process  (cf.  p.  549), 


CARBON  AND  THE  OXIDES  OF  CARBON  569 

the  production  in  the  U.  S.  being  2542  short  tons  (1915).  A  power- 
ful alternating  current  is  passed  through  a  mass  of  granular  anthra- 
cite, mixed  with  pitch  and  a  little  sand  (Acheson's  process).  The 
mixture  (3  tons)  is  piled  between 
the  electrodes  (Fig.  127)  and,  on 
account  of  its  high  resistance, 
becomes  strongly  heated.  The 
change  occupies  24-30  hours. 

Graphite  is  now  used  exclu- 
sively for  making  the  anodes  in 
the  electrolytic  manufacture  of 
chlorine  and  in  related  processes.  FlG  127. 

Mixed  with  fine  clay  it  forms  the 

"lead"  of  lead  pencils,  first  used  in  the  sixteenth  century.*  Mixed 
with  clay  it  is  used  also  for  making  crucibles,  which  withstand  high 
temperatures  and  serve  for  melting  and  casting  steel  and  high  melt- 
ing alloys.  As  " black-lead"  it  forms  stove  polish,  the  layer  of  fine 
scales  protecting  the  iron  against  rusting.  It  is  employed  as  a 
lubricant  in  cases  where  oil  would  be  decomposed  by  the  heat  and 
where  wooden  surfaces  are  in  contact. 

Chemical  Properties  of  Carbon.  - —  Diamond,  graphite,  and 
amorphous  carbon  probably  differ  from  one  another,  not  merely  in 
physical  properties,  but  also  chemically.  Certainly  the  stability  of 
compounds  containing  many  units  of  carbon  in  their  molecules  indi- 
cates a  great  tendency  of  carbon  to  combine  with  itself,  and  gives 
plausibility  to  the  belief  that  the  molecule  of  free  carbon  may  itself  be 
complex.  Differences  in  the  arrangement  of  the  atoms  (p.  471) 
account  for  the  variety  in  the  crystalline  forms  of  the  element. 
Amorphous  carbon  is  the  least  stable  of  the  three,  for  it  liberates  most 
heat  in  entering  into  combination.  Since  graphite  is  formed  at  high 
temperatures,  and  diamonds  turn  into  a  black  mass  under  the  same 
conditions,  we  may  presume  that  graphite  is  the  most  stable,  at  least 
at  3000°. 

The  most  common  uses  of  carbon  depend  upon  its  great  tendency 
to  unite  with  oxygen,  forming  carbon  dioxide  CO2.  Under  some 
circumstances  carbon  monoxide  CO  (see  below)  is  produced.  Aside 
from  the  direct  employment  of  this  action  for  the  sake  of  the  heat 
which  is  liberated,  it  is  used  also  in  the  reduction  of  ores  of  iron, 

*  Priestley  was  the  first  to  suggest  the  use  of  caoutchouc  (raw  rubber)  as 
an  eraser. 


570  INORGANIC   CHEMISTRY 

copper,  zinc,  and  many  other  metals.  When,  for  example,  finely 
powdered  cupric  oxide  and  carbon  are  heated,  copper  is  obtained. 
The  gas  given  off  is  either  carbon  dioxide,  or  a  mixture  of  this  with 
carbon  monoxide,  according  to  the  proportion  of  carbon  used : 

CuO  +  C-+    Cu  +  CO, 
2CuO  +  C  ->  2Cu  +  CO2. 

The  union  with  hydrogen  is  ordinarily  too  slow  to  be  observed. 
But  when  the  carbon  is  mixed  with  pulverized  nickel  (contact  agent) , 
and  hydrogen  is  passed  over  the  mixture  at  250°,  methane  CH4 
is  formed  (99  per  cent).  The  action  is  reversible  and  exothermal, 
and  is  therefore,  at  higher  temperatures,  less  complete  (cf.  p.  305), 
at  850°  reaching  only  1.5  per  cent.  On  the  other  hand,  an  electric 
arc,  between  carbon  poles  in  an  atmosphere  of  hydrogen,  gives 
traces  of  acetylene  C2H2  (q.v.)  this  action  being  endothermal.  The 
other  compounds  of  carbon  and  hydrogen  are  all  obtained  by  indirect 
reactions. 

At  the  high  temperatures  produced  in  the  electric  furnace,  carbon 
unites  with  many  metals  and  some  non-metals.  Compounds  formed 
in  this  way  are  known  as  carbides,  such  as  aluminium  carbide  A14C3, 
calcium  carbide  CaC2,  and  carborundum  CSi  (see  below). 

Carbon  Disulphide  CS2.  —  This  compound  is  made  by  direct 
union  of  sulphur  vapor  and  glowing  charcoal.  An  electric  furnace 
like  that  in  Fig.  125  (p.  548)  is  employed.  The  substance  comes  off 
as  a  vapor  and  is  condensed. 

Carbon  disulphide  is  a  colorless,  highly  refracting  liquid  (b.-p. 
46°).  Traces  of  other  compounds  give  the  commercial  article  a 
disagreeable  smell.  It  burns  in  air,  forming  carbon  dioxide  and 
sulphur  dioxide.  It  is  an  important  solvent  for  sulphur  and  caout- 
chouc (rubber),  and  dissolves  iodine  and  phosphorus  freely.  Large 
quantities  are  employed  also  in  the  destruction  of  prairie  dogs  and 
ants,  and  for  freeing  grain  elevators  of  rats  and  mice. 

Carbon  Tetrachloride  CCh.  —  This  compound  is  manufac- 
tured by  leading  dry  chlorine  into  carbon  disulphide  containing 
a  little  iodine  (contact  agent)  in  solution: 

CS2  +  3C12  ->  CC14  +  S2C12. 

The  carbon  tetrachloride  (b.-p.  77°)  is  first  distilled  off,  and  the 
sulphur  monochloride  (b.-p.  136°)  is  purified  for  use  in  vulcanizing 
rubber. 


CARBON  AND  THE  OXIDES  OF  CARBON 


571 


Carbon  tetrachloride  is  a  colorless  liquid  which  dissolves  fats, 
tars,  and  many  other  organic  compounds.  It  is  used  to  take  the 
oil  or  grease  out  of  wool,  linen,  oil-bearing  seeds  and  bones.  It  has 
a  great  advantage  over  gasoline  (petrol)  and  benzine  (see  p.  586), 
which  can  be  used  for  similar  purposes,  in  that  it  is  non-inflammable. 
"Carbona,"  used  for  removing  stains  from  clothing,  gloves,  and 
shoes,  is  benzine  to  which  sufficient  carbon  tetrachloride  has  been 
added  to  render  the  mixture  non-inflammable.  "Pyrene"  fire 
extinguishers  contain,  mainly,  carbon  tetrachloride.  The  tem- 
perature of  the  burning  material  is  lowered,  because  heat  is  con- 
sumed in  vaporizing  the  liquid  and,  at  the  same  time,  the  vapor 
displaces  the  air  and  stops  the  combustion. 

Calcium  Carbide  CaC%  and  Carborundum  SiC.  —  Calcium 
carbide  is  manufactured  in  an  electric  furnace  (Thomas  Willson, 
a  Canadian)  by  the  interaction  of  finely  pulverized  limestone  or 
quicklime  with  coke: 

CaO  +  3C  ->  CaC2  +  CO. 

The  operation  is  a  continuous  one,  the  materials  being  thrown  into  the 
left  side  of  the  drum  (Fig.  128,  diagrammatic),  and  the  product  re- 
moved on  the  right.  The  car- 
bon poles  are  fixed.  The  arc 
having  been  established,  the 
drum  is  rotated  slowly  as  the 
carbide  accumulates.  The  cur- 
rent enters  by  one  carbon,  passes 
through  the  carbide,  and  leaves 
by  the  other.  The  high  resist- 
ance of  the  partially  trans- 
formed material  causes  the  pro- 
duction of  the  heat.  When  the 
action  in  one  layer  approaches 
completion,  the  resistance  falls, 
the  current  increases,  and  an 
armature  round  which  the  wire 
passes  (not  shown  in  Fig.  128)  FIQ.  128. 

comes  into  operation  and  turns 

the  drum.  In  this  way  the  carbide  just  formed  is  continuously 
moved  away  from  the  carbons,  and  new  material,  introduced  on 
the  left,  falls  into  the  path  of  the  current.  The  iron  plates  which 


572  INORGANIC  CHEMISTRY 

form  the  circumference  of  the  drum  are  added  on  the  left  and  re- 
moved on  the  right,  where  also  the  carbide  is  broken  out  with  a 
chisel.  The  drum  revolves  once  in  about  three  days.  The  product 
is  used  for  making  acetylene  (q.v.). 

Carborundum,  or  carbide  of  silicon  SiC,  of  which  hundreds  of 
tons  are  manufactured  annually  at  Niagara  Falls  (Acheson's  process), 
is  made  in  an  electric  furnace  of  the  type  shown  in  Fig.  127  (p.  569). 
A  mixture  of  coke  and  sand  (silicon  dioxide  Si02),  containing  some 
sawdust,  is  piled  between  the  terminals,  with  a  core  of  granular 
carbon  to  carry  most  of  the  current.  The  resistance  produces  a  high 
temperature  (1950°),  and  the  sand  is  reduced: 

3C  +  Si02  ->  SiC  +  2CO. 

The  carborundum  remains,  often  in  beautifully  crystalline  form. 
It  is  exceedingly  hard  (Appendix  II),  and  after  pulverization  and 
mixing  with  a  filler,  is  molded  into  grinding  wheels  and  whet-stones. 
Carborundum  decomposes  at  2220°.  It  is  not  affected  by  water 
or  acids,  but  is  decomposed  by  alkalies. 

The  Oxides  of  Carbon.  —  Four  oxides  of  carbon  are  known, 
of  which  two,  namely  the  dioxide  C02  and  the  monoxide  CO,  are 
familiar.  Carbon  suboxide  C3O2,  and  mellitic  anhydride  Ci2Og  (50 
per  cent  carbon  and  50  per  cent  oxygen)  are  best  classed  as  organic 
compounds.  Two  other  oxides,  C5O5  and  C6O6,  are  known  only  in 
the  hydrated  forms,  namely  leuconic  acid  and  triquinoyl,  respectively. 

CARBON  DIOXIDE  AND  CARBONIC  ACID 

Occurrence.  —  Carbon  dioxide  is  present  in  the  atmosphere, 
and  issues  from  the  ground  in  large  quantities  in  certain  neighbor- 
hoods, as,  for  example,  near  the  Lake  of  Laach,  in  the  so-called 
Valley  of  Death  in  Java,  and  in  the  Grotta  del  Cane  near  Naples. 
Effervescent  mineral  waters  contain  it  in  solution,  and  their  effer- 
vescence is  caused  by  the  escape  of  the  gas  when  the  pressure 
is  reduced.  Well-known  waters  of  this  kind  are  those  of  Selters 
(whence,  by  a  singular  perversion,  the  English  word  seltzer  is  de- 
rived), of  Vichy  and  of  the  Geyser  Spring  at  Saratoga. 

Modes  of  Formation.  —  1.   Carbon    dioxide  is  produced  by 

combustion  of  carbon  with  an  excess  of  oxygen :  C  +  O2  — >  CO2. 
The  combustion  of  all  compounds  of  carbon,  as  well  as  the  slow  oxida- 


CARBON  AND  THE  OXIDES  OF  CARBON  573 

tion  in  the  tissues  of  plants  and  animals,  leads  to  the  formation  of  the 
same  substance.  The  product  from  burning  carbon  is  naturally 
mixed  with  at  least  four  times  its  volume  of  atmospheric  nitrogen. 
To  secure  carbon  dioxide  for  commercial  purposes  from  this  source, 
the  gas  is  led  under  pressure  into  a  solution  of  potassium  carbonate, 
which  absorbs  the  carbon  dioxide: 

C02  (gas)  *=±  CO2  (dslvd)  +  H20  <=±  H2CO3  +  K2C03  <=±  2KHCO3. 

When  the  pressure  is  reduced  by  a  pump,  all  the  actions  are  reversed, 
and  the  gas  escapes  in  pure  form.  The  same  solution,  with  occasional 
purification,  can  be  used  an  indefinite  number  of  times. 

2.  It  was  Joseph  Black  (1757)  who  first  recognized  the  gas  as  a 
distinct  substance.     He  observed  its  formation  when  marble  or  mag- 
nesium carbonate  was  heated: 

CaC03  <=*  CaO  +  CO2, 

and  named  the  gas  "  fixed  air"  from  the  fact  that  it  was  contained  in 
these  solids.  The  above  action  had  been  used  for  centuries  in 
making  quicklime  (calcium  oxide).  All  common  carbonates,  except- 
ing the  normal  carbonates  of  potassium  and  sodium,  decompose  in 
this  way,  leaving  the  oxide  of  the  metal  or  the  metal  itself  (p.  130). 

3.  Black  found  that  the  gas  was  also  produced  when  acids  acted 
upon  carbonates,  and  this  method  is  commonly  employed  in  the 
laboratory  : 


CaCO3  (solid)^CaC03  (dslvd)  ^Ca+++C03=K   TT  rn  <_„  n_Lrn 
2HC1  (ddvd)«2Cr+2H+    I" 

Since  the  carbonic  acid  is  very  slightly  ionized,  the  action  is  like 
that  of  acids  on  sulphites  (p.  424).  Since,  however,  the  carbonate 
of  calcium  (marble)  is  very  slightly  soluble,  so  that  an  additional 
equilibrium  controls  its  solution,  the  action  is  like  that  of  acids  on 
ferrous  sulphide  (p.  419).  The  apparatus  shown  in  Fig.  41  (p.  119) 
is  used. 

4.  Carbon  dioxide  is  also  a  product  of  the  fermentation  of  sugar 
(q.v.),  as  Black  had  the  credit  of  showing.  It  is  formed  also,  with  the 
assistance  of  bacteria,  in  the  decay  of  animal  and  vegetable  matter 
(p.  91). 

Physical  Properties.  —  Carbon  dioxide  is  a  colorless,  odorless 
gas.  It  is  one-half  heavier  than  air.  The  G.M.V.  weighs  44.26  g. 
The  critical  temperature  is  31.35°.  The  solid  melts  at  —56°,  having 


574  INORGANIC  CHEMISTRY 

a  vapor  pressure  of  5.3  atmospheres.  The  solid  has  a  vapor  pressure 
of  1  atmos.  at  -79°.  The  density  of  the  liquid  at  0°  is  0.95.  At  0° 
its  vapor  tension  is  35.4  atmospheres  and  at  20°  59  atmospheres. 
It  must  be  preserved,  therefore,  in  very  strong  cylinders  of  mild 
steel.  Large  quantities  of  it,  often  collected  from  fermentation  vats, 
are  sold  in  such  cylinders,  and  used  in  operating  beer-pumps  and  in 
making  aerated  waters.  When  the  liquid  is  allowed  to  flow  out 
into  an  open  vessel  or,  better  still,  into  a  cloth  bag  (nonconductor 
of  heat),  it  cools  itself  by  its  own  evaporation  and  forms  a  white, 
snowlike  mass.  Solid  carbon  dioxide  evaporates  at  —79°  without 
melting,  since  at  that  temperature  it  exercises  1  atmosphere  pressure, 
and  the  heat  from  the  surroundings  is  used  as  heat  of  vaporization 
instead  of  being  employed  in  raising  the  temperature  to  the  melting- 
point  (-56°). 

The  solid  is  used  in  the  laboratory  as  a  cooling  agent,  being  often 
mixed  with  ether  to  give  closer  contact  with  the  vessel  (  —  80°). 
Mercury  (m.-p.  —40°)  is  easily  frozen  by  the  mixture. 

The  great  contrast  in  the  speeds  of  a  chemical  change  at  two 
temperatures  (cf.  p.  93)  may  be  illustrated  by  putting  a  minute  piece 
of  sodium  in  some  30  per  cent  hydrochloric  acid  which  has  been 
cooled  in  the  above  mixture.  Hardly  any  interaction  can  be  ob- 
served. But  if  the  temperature  of  the  acid  is  allowed  to  rise,  the 
action  becomes  more  and  more  rapid,  and  ends  by  being  explosively 
violent. 

Carbon  dioxide  gas  (760  mm.  and  15°)  dissolves  in  its  own  volume 
of  water.  Up  to  four  or  five  atmospheres  Henry's  law  (p.  188) 
describes  its  solubility  accurately.  An  aqueous  solution,  prepared 
under  a  pressure  of  8-10  atmospheres,  is  familiarly  known  as  soda 
water,  or  carbonated  water. 

Chemical  Properties.  —  Carbon  dioxide  is  a  stable  compound. 
At  2000°  (760  mm.  press.)  the  dissociation  reaches  1.8  per  cent 
(2200°,  4.9  per  cent;  2500°,  15.8  per  cent):  2C02  ^  2CO  +  O2,  or 
about  the  same  as  that  of  water. 

The  more  active  metals,  like  magnesium,  burn  brilliantly  when 
ignited  in  a  hollow  lump  of  solid  carbon  dioxide,  producing  the  oxide 
of  the  metal  and  free  carbon.  Less  active  metals,  such  as  zinc  and 
iron,  when  heated  in  a  stream  of  the  gas,  give  an  oxide  of  the  metal  and 
carbon  monoxide  (q.v.). 

Carbon  dioxide  unites  directly  with  many  oxides,  particularly 
those  of  the  more  active  metals,  such  as  the  oxides  of  potassium, 


CARBON  AND  THE  OXIDES  OF  CARBON  575 

sodium,  calcium,  etc.,  giving  the  carbonates.  Hence  the  decomposi- 
tion of  calcium  carbonate  by  heating  (p.  573)  is  a  reversible  action, 
which  proceeds  in  the  opposite  direction  when  a  sufficient  pressure  of 
carbon  dioxide  is  used  (cf.  p.  153). 

Carbon  dioxide,  when  dissolved  in  water,  forms  an  unstable  acid  : 

H  O  H  -  0 

H2O  +  CO2  «=±  H2CO3,     or          \  o  +  cf        ->  XC  =  O. 


The  name  carbonic  acid  is  frequently,  though  improperly,  given  to 
the  anhydride  CO2,  which  has  no  acid  properties. 

Chemical  Properties  of  Carbonic  Acid  lf2C03.  —  The  solu- 
tion of  carbon  dioxide  in  water  exhibits  the  properties  of  a  weak  acid. 
It  conducts  electricity,  although  not  well.  It  turns  litmus  red, 
though  not  so  decidedly  as  do  strong  acids.  Its  feebleness  is  due, 
however,  not  exclusively  to  the  small  degree  of  ionization,  but  also 
to  the  fact  that  ordinary  solutions  of  carbon  dioxide  are  necessarily 
very  dilute.  The  ionization  takes  place  chiefly  according  to  the 
equation  : 

LH2C03  <=±  H+  +  HCOr. 

In  a  deci-normal  solution,  less  than  two  molecules  of  the  acid  in  a 
thousand  are  ionized.  The  conditions  of  equilibrium  between  the  gaa 
and  the  solution  are  precisely  similar  to  those  described  under  sul- 
phurous acid  (p.  444). 

Carbonates  and  Bicarbonates.  —  When  excess  of  an  aqueous 
solution  of  carbonic  acid  is  mixed  with  a  solution  of  a  base  like 
sodium  hydroxide,  or,  as  the  operation  is  more  usually  performed, 
when  carbon  dioxide  is  passed  directly  into  a  solution  of  the  alkali, 
water  is  formed  and  the  acid  carbonate  (bicarbonate)  of  sodium 
remains  dissolved: 

H2C03  +  NaOH  *=>  H20  +  NaHCO3,     or     H+  +  OH~  ^  H2O. 

Although  the  bicarbonate  is  technically  an  acid  salt,  its  solution  is 
neutral  on  account  of  the  exceedingly  slight  dissociation  of  the  HCO3~ 
ion.  By  addition  of  an  equivalent  of  sodium  hydroxide  to  the 
solution  of  the  bicarbonate  the  normal  carbonate  is  obtained  : 

NaOH  +  NaHCO3^±H2O  +NasC03,  or  OH 


576  INORGANIC  CHEMISTRY 

This  solution,  like  that  of  all  salts  of  a  strong  base  and  a  feeble  acid 
(of.  p.  399),  is  alkaline  in  reaction.  This  is  because  the  tendency  to 
form  the  very  slightly  ionized  HCO3~  makes  the  foregoing  ionic 
action  noticeably  reversible  (cf.  pp.  419,  559). 

The  normal  carbonates,  with  the  exception  of  those  of  potassium, 
sodium,  and  ammonium,  are  insoluble  in  water,  and  may  be  obtained 
by  precipitation  when  the  proper  ions -are  employed.  For  example: 

BaCl2  +  Na2CO3  +±  BaC03  j  +  2NaCl,  or  Ba++  +  CO-r  +±  BaC03| . 

The  aqueous  solution  of  carbon  dioxide  interacts  with  solutions 
of  barium  and  calcium  hydroxides  in  a  similar  manner: 

Ca(OH)2  +  H2C03  fc?  CaCOsi  +  2H20. 

These  precipitations  are  used  as  tests  for  carbon  dioxide  and  as  a 
means  of  estimating  its  amount  in  a  sample  of  air  (p.  501). 

Excess  of  carbon  dioxide  converts  calcium  carbonate  into  the  more 
soluble  bicarbonate,  and  hence  considerable  quantities  of  "lime" 
(hardness,  q.v.)  are  frequently  held  in  solution  by  natural  waters, 
all  of  which  contain  carbon  dioxide  in  solution: 

H2C03  +  CaC03  <=*  Ca(HC03)2. 

A  considerable  excess  of  carbon  dioxide  is  required  to  convert  the 
whole  of  the  carbonate  into  the  soluble  bicarbonate,  since  the  action 
is  markedly  reversible.  In  the  same  fashion,  the  carbonates  of  iron 
FeC03  (chalybeate  water),  magnesium,  and  zinc  are  dissolved  as 
bicarbonates  in  water  containing  free  carbonic  acid.  In  fact,  'the 
solution,  transportation,  and  deposition  of  all  these  carbonates  take 
place  in  nature  on  a  large  scale  by  the  alternate  progress  and  re- 
versal of  this  action. 

Uses  of  Carbon  Dioxide.  —  The  employment  of  the  gas  for 
impregnating  aerated  waters  has  been  mentioned.  The  gas  is  used 
in  immense  quantities  in  the  manufacture  of  sodium  bicarbonate 
NaHCO3  (baking  soda),  of  sodium  carbonate  Na2CO3,10H20  (wash- 
ing soda),  and  of  white  lead,  a  basic  carbonate  of  lead  Pb3(OH)2(C03)2. 

Since  carbon  dioxide  is  already  fully  oxidized,  it  does  not  burn, 
and  since  it  is  very  stable,  ordinary  combustibles  will  not  burn  in 
it.  A  small  percentage  of  it  will  destroy  the  power  of  air  to  sup- 
port combustion.  For  this  reason,  portable  fire  extinguishers 
contain  a  dilute  solution  of  sodium  bicarbonate,  and  a  bottle  of 


CARBON  AND  THE  OXIDES  OF  CARBON  577 

sulphuric  acid.  When  the  tank  is  inverted,  the  acid  flows  into  the 
solution: 

2NaHC03  +  H2SO4  ^  Na*S04  +  2H2C03  *±  2H2O  +  2CO2. 

The  liquid  is  saturated  with  the  gas  and  the  excess,  rising  to  the 
top,  by  its  pressure  forces  the  solution  out  through  the  nozzle.  The 
liquid  is  more  effective  than  an  equal  amount  of  water,  because 
the  carbon  dioxide  it  carries  mixes  with  the  surrounding  air. 

The  most  wonderful  chemical  change  which  carbon  dioxide 
undergoes  is  perhaps  the  most  useful  to  mankind,  and  at  the  same 
time  the  one  least  understood.  This  is  the  action  by  which  plants 
use  it  as  food  (see  p.  579). 

CARBON  MONOXIDE  CO 

Preparation.  —  In  the  laboratory,  carbon  monoxide  may  be 
obtained  by  heating  oxalic  acid,  a  solid,  white,  crystalline  substance, 
in  a  flask  with  concentrated  sulphuric  acid.  The  latter  is  here 
employed  simply  as  a  dehydrating  agent  (p.  438) : 

H2C204  ->  C02  +  CO  +  H20. 

To  obtain  pure  carbon  monoxide  from  this  mixture,  it  is  necessary 
to  remove  the  carbon  dioxide,  by  passing  the  gas  through  a  solu- 
tion of  potassium  hydroxide  contained  in  a  wash  bottle.  By  using 
formic  acid,  or  sodium  formate,  with  sulphuric  acid,  the  presence 
of  the  carbon  dioxide  is  avoided:  HCHO2  ->  CO  +  H2O. 

We  commonly  observe  the  blue  flame  of  burning  carbon  mon- 
oxide playing  on  the  surface  of  a  coal  fire.  The  gas  is  produced  by 
the  passage  of  the  carbon  dioxide,  which  is  first  formed,  through 
the  upper  layers  of  heated  coal:  C02  +  C  — » 2CO.  A  similar 
reduction  of  carbon  dioxide  is  produced  when  the  gas  is  led  over  a 
metal,  such  as  zinc,  and  heat  is  applied :  CO2  +  Zn  — >  ZnO  +  CO. 

Producer  Gas  and  Water  Gas.  —  When  coke  and  air  are  used 
in  the  reaction  mentioned  above,  the  mixture  of  carbon  monoxide 
(about  33  per  cent)  and  nitrogen  (about  66  per  cent)  obtained  is 
called  producer  gas.  It  is  combustible  and  is  used  in  factories  for 
heating  and  to  drive  gas  engines  for  power. 

When  steam  is  driven  through  white  hot  coke  or  anthracite,  a 
mixture  of  hydrogen  and  carbon  monoxide,  known  as  water  gas 
is  produced: 

C  4-  H20  ->  CO  +  H2  -  28,300  cal. 


578  INORGANIC  CHEMISTRY 

The  coke,  piled  in  a  brick-lined,  cylindrical  structure,  is  brought 
to  vigorous  combustion  by  blowing  in  air  for  ten  minutes.  Then 
steam  is  substituted  for  the  air.  Since  the  interaction  takes  place 
with  absorption  of  heat  (is  endothermal,  see  equation),  in  about 
five  minutes  the  coke  becomes  too  cool.  Air  is  then  substituted 
for  steam,  and  so  on  alternately.  The  gas  is  collected  while  the 
steam  is  turned  on,  and  contains  equal  volumes  of  the  two  gases, 
together  with  some  carbon  dioxide  (4-7  per  cent),  nitrogen  (4-5 
per  cent)  and  oxygen  (1  per  cent).  The  gas  is,  therefore,  almost 
wholly  combustible  and  is  used  as  a  source  of  heat,  and  for  driving 
gas  engines  to  furnish  power.  It  is  used  also  for  making  illuminat- 
ing gas  (q.v.).  Since  carbon  monoxide  is  more  easily  liquefied  than 
is  hydrogen,  the  latter  gas  is  obtained,  for  commercial  use,  by  passing 
water  gas  through  a  liquefier. 

When  both  steam  and  air  are  driven  together  over  burning  coke, 
the  latter  is  able  to  burn  continuously,  and  a  fuel  gas  which  is  a 
cross  between  producer  gas  and  water  gas  is  obtained. 

Fuel  gases  are  used  on  a  large  scale  in  steel  works,  and  other 
factories.  They  give  a  uniform  and  easily  regulated  heat,  they 
leave  no  ash,  and  their  use  involves  no  labor  for  stoking.  As  gases, 
also,  they  can  be  used  in  structures  in  which  coal,  as  a  solid,  could 
not  be  employed. 

Physical  Properties.  —  Carbon  monoxide  is  a  colorless  gas, 
with  a  metallic  taste  and  odor  (poisonous!).  It  is  very  slightly 
soluble  in  water.  Its  density  is  almost  the  same  as  that  of  air,  for  the 
G.M.V.  weighs  28  g.  When  liquefied  it  boils  at  -  190°. 

Chemical  Properties.  —  All  the  chemical  properties  of  carbon 
monoxide  are  referable  to  the  fact  that  in  it  the  element  carbon 
is  bivalent:  C  =  O.  The  compound  is  in  fact  unsaturated,  and 
combines  with  oxygen,  chlorine,  and  other  substances  directly. 
Thus  the  gas  burns  in  the  air,  uniting  with  oxygen  to  form  carbon 
dioxide.  Again,  iron  (q.v.)  is  manufactured  by  the  reduction  of  the 
oxide  of  iron  by  gaseous  carbon  monoxide  in  the  blast  furnace: 

SCO  <=±  2Fe  +  3C02. 


In  sunlight  carbon  monoxide  unites  directly  with  chlorine  to  form 
carbonyl  chloride  COC12.  It  is  absorbed  by  a  solution  of  cuprous 
chloride  in  hydrochloric  acid,  forming  a  compound  said  to  be 
CuCOCl,H20.  It  unites  directly  with  certain  metals,  notably 


CARBON  AND  THE  OXIDES  OF  CARBON  579 

nickel  and  iron,  with  which  it  forms  the  so-called  nickel  carbonyl 
Ni(CO)4  (q>v.)  and  iron  carbonyl,  respectively.  The  gas  reduces 
Fehling's  solution  (q.v.),  and  precipitates  silver  from  ammonio-silver 
nitrate  solution. 

The  gas  is  an  active  poison.  When  inhaled  it  unites  with  the 
haemoglobin  of  the  blood,  to  the  exclusion  of  the  oxygen  which 
forms  with  the  haemoglobin  a  less  stable  compound  (cf.  p.  92).  A 
quantity  equivalent  to  about  10  c.c.  of  the  gas  per  kilo,  weight  of  the 
animal  is  sufficient  to  produce  death,  about  one-third  of  the  whole 
haemoglobin  having  entered  permanently  into  combination  with 
carbon  monoxide.  One  volume  in  800  volumes  of  air  produces 
death  in  about  thirty  minutes.  This  gas  is  the  chief  poisonous 
substance  in  illuminating  gas.  The  poisonous  effect  of  tobacco 
smoke,  when  inhaled,  is  partly  due  to  the  carbon  monoxide  pro- 
duced by  incomplete  combustion.  Nicotine,  although  contained  in 
tobacco  leaves,  is  unstable,  and  is  decomposed  by  the  heat.  Traces 
of  other  irritant  organic  compounds,  however,  are  contained  in  the 
smoke. 

The  quantities  of  heat  given  out  by  the  successive  unions  of  two  units  of  oxy- 
gen with  one  unit  of  amorphous  carbon  are  worth  recording: 

C  +  O-» CO  +  29,650 calories,    and    CO  +  O -» CO2  +  68,000 calories. 

It  will  be  seen  that  the  addition  of  the  second  atom  of  oxygen  appears  to  cause  the 
evolution  of  a  very  much  larger  amount  of  heat  than  does  that  of  the  first.  It 
must  be  remembered,  however,  that  the  carbon  monoxide  is  gaseous,  while  the 
carbon  in  the  first  equation  is  solid.  The  heats  produced  by  the  unions  of  the 
two  units  are  probably  not  very  different,  but  in  the  first  case  a  large  amount  of 
the  heat  is  used  up  in  bringing  the  carbon  into  the  gaseous  condition. 

Carbon  Suboxide  C3O2.  —  This  oxide  is  obtained  by  the  action 

of  phosphorus  pentoxide  (dehydrant)  on  malonic  acid.  H2(COfe)2CH2 
— >  2H2O  +  C3O2 1 .  It  is  a  colorless  liquid,  boiling  at  7°.  The  vapor 
has  an  unpleasant  odor.  With  water  it  gives  malonic  acid,  of  which 
it  is  the  anhydride.  If  kept  for  a  day,  it  changes  into  a  dark  red 
solid  of  the  same  composition. 

Carbon  Dioxide  as  Plant  Food.  —  The  walls  of  the  cells  which 

form  the  framework  of  a  plant  are  made  of  cellulose  (C«HioO5)x. 
In  the  cells,  especially  those  in  certain  parts  of  the  plant,  granules 
of  starch  (CeHuA^  are  found,  and  in  the  fruit,  sugars  CeH^Oe  or 
stored.  The  plant  contains  also  proteins.  These  sub- 


580  INORGANIC  CHEMISTRY 

stances  contain  carbon,  hydrogen,  oxygen,  nitrogen,  sulphur,  and  phos- 
phorus, and  plant  food  must  furnish  these  elements.  Compounds  of 
potassium  are  also  required.  Hence,  in  addition  to  large  amounts 
of  water  ascending  through  the  roots  and  stem,  carrying  sufficient 
quantities  of  soluble  compounds  of  the  four  elements  last  named,  all 
plants  require  an  abundant  supply  of  carbon  in  absorbable  form. 
Now,  this  comes  from  atmospheric  carbon  dioxide,  admitted  through 
minute  openings  (stomata)  situated  mainly  on  the  lower  surfaces  of 
the  leaves.  Comparison  of  the  formulae  C02  and  CeHioOs,  shows  at 
once  that  the  assimilation  of  the  carbon  dioxide  of  the  plant  must 
involve  reduction.  The  chlorophyll  (green  matter)  and  protoplasm 
in  the  leaves  act  upon  the  carbon  dioxide,  causing  oxygen  gas  to  be 
liberated.  It  is  believed  that  the  carbon  dioxide  is  reduced  to  formal- 
dehyde CH2O,  and  from  this  sugars  can  be  made  in  the  laboratory: 
6CH20  — >  CeH^Oe.  It  is  supposed  that  the  various  chemical  com- 
pounds which  plants  construct  in  large  quantities,  such  as  sugar, 
starch,  and  cellulose,  are  built  up  as  the  result  of  actions  like  this. 
In  a  rough  fashion,  and  disregarding  the  steps  by  which  the  process 
takes  place,  we  may  represent  the  chemical  change  by  means  of  the 
thermochemical  equation: 

6C02  +  5H20  +  671,000  cal.  ->  C6Hi005  +  602. 

These  figures  indicate  roughly  (p.  35)  the  amount  of  energy  stored  for 
cellulose,  and  the  values  for  the  other  compounds  are  of  the  same 
order.  This  action  goes  on  only  in  sunlight  (see  next  section)  and  if 
green  leaves  are  placed  under  water  saturated  with  carbon  dioxide, 
oxygen  is  given  off  and  can  be  collected. 

The  enormous  amount  of  energy  absorbed  in  the  action,  and 
represented  in  terms  of  heat  in  the  equation,  is  furnished  by  the 
sunlight.  It  may  be  added  that  plants,  like  animals,  also  use  some 
oxygen  and  produce  some  carbon  dioxide,  but  this  process  is  entirely 
overborne  in  daylight,  and  is  noticeable  only  in  the  dark. 

-The  energy  that  does  the  world's  work  comes  mainly  from  two 
sources,  namely,  water  power  and  the  combustion  of  wood  or  coal 
(which  is  fossil  wood).  The  water  comes  from  vapor,  generated 
by  the  sun's  heat,  condensed  as  rain,  and  ultimately  feeding  the 
rivers.  The  source  of  the  energy  in  wood  and  coal  is  now  apparent. 
When  wood,  which  is  largely  cellulose  (CeHioOs)*,  burns,  it  gives 
carbon  dioxide,  water,  and  heat.  In  fact,  its  combustion  is  rep- 
resented by  the  above  equation,  when  read  backwards.  Thus,  the 
sunlight,  through  the  machinery  of  the  plant,  takes  carbon  dioxide 


CARBON  AND  THE  OXIDES  OF  CARBON       581 

and  water,  supplies  the  energy  (as  light),  and  gives  us  wood  and 
oxygen.  And  the  wood  and  oxygen,  when  burned,  give  us  back 
the  original  substances,  and  the  equivalent  of  the  original  energy  in 
the  form  of  heat.  Thus,  the  two  sources  of  energy  turn  out  to  be 
the  same,  namely  the  sun's  rays. 

If,  instead  of  burning  the  starch  of  the  plant,  we  consume  it  as 
food,  it  goes  through  several  changes  instead  of  one.  But  the  final 
products  are  the  same,  namely  carbon  dioxide  and  moisture,  given 
off  through  our  lungs  and  skin,  and  heat  and  other  forms  of  energy 
such  as  are  developed  in  animals.  Thus,  whether  we  use  our  muscles, 
a  steam  engine,  or  a  water  turbine  to  do  work,  sunlight  is  in  each 
case  the  ultimate  source  of  the  energy  employed. 

It  should  be  noted  that  the  energy  is  not  stored  exclusively  in  the  coal,  but 
is  shared  between  coal  and  the  oxygen  of  the  air.  If  our  atmosphere  consisted 
of  compounds  of  carbon,  then  the  material  corresponding  to  stores  of  coal  would 
have  to  be  oxygen  or  compounds  of  oxygen,  and  we  should  be  likely  then  to 
speak  of  the  energy  as  being  stored  for  us  and  sold  in  the  form  of  oxygen.  That 
we  are  in  the  habit  of  speaking  of  it,  at  present,  as  going  with  the  carbon  is  because 
the  oxygen  of  the  air  is  supplied  free  of  charge,  while  the  coal  and  wood  have 
to  be  purchased. 

Photochemical  Action.  —  We  have  seen  that  light  may  simply 
act  catalytically,  as  on  a  mixture  of  hydrogen  and  chlorine  (p.  222)  or 
an  aqueous  solution  of  hypochlorous  acid  (p.  474).  These  actions  in- 
volve the  liberation  of  energy  and  go  on  spontaneously  (cf.  p.  35) 
under  proper  conditions.  On  the  other  hand,  light  may  actually  be 
consumed  in  large  amount  in  producing  a  chemical  change,  as  in 
decomposing  silver  chloride  (p.  19),  or  in  the  above  instance.  All 
wave  lengths  of  light,  which  is  the  same  as  to  say  all  colors  of  light, 
are  not  equally  active  in  any  one  case.  But  there  is  no  particular  set 
of  wave  lengths  which  is  of  special  chemical  activity.  In  the  action 
on  silver  chloride,  green  and  blue  light  is  very  active,  while  red  is 
almost  without  effect.  Here,  in  the  actions  in  which  chlorophyll  is 
concerned,  it  is  the  red  and  yellow  light  that  produces  the  chemical 
change,  and  a  plant  exposed  to  blue  light  (e.g.,  by  shading  with  blue 
glass)  will  assimilate  none  of  the  carbon  dioxide  in  the  air  surrounding 
it.  The  chemical  substances  in  the  retina  of  the  eye  seem  to  resemble 
those  in  the  leaves  of  plants,  for  they  are  most  affected  by  red  and 
yellow  light.  To  put  this  another  way,  a  spectrum  of  uniform 
intensity  throughout,  when  viewed  by  a  plant  or  a  human  eye,  would 
appear  to  be  brightest  in  the  red  and  yellow  portions,  while  a  con- 


582  INORGANIC   CHEMISTRY 

siderable  stretch  towards  the  blue  extremity  would  actually  be 
invisible.  On  the  other  hand,  to  an  eye  in  which  the  active  substance 
was  silver  chloride,  if  such  an  eye  could  be  imagined  as  existing,  the 
red  end  would  be  invisible  and  the  blue  and  ultra-violet  would  be  the 
most  brilliant  parts. 

CARBONYL  CHLORIDE  AND  UREA 

Carbonyl  Chloride  COClz.  —  This  substance  is  also  named 
phosgene  (Gk.  <£<<%  light;  ytvvavt  to  produce),  on  account  of  its  forma- 
tion by  the  catalytic  influence  of  sunlight  (p.  581).  On  a  commercial 
scale  it  is  obtained  by  passing  the  mixed  carbon  monoxide  and 
chlorine  over  animal  charcoal  (contact  agent).  It  is  a  liquid  which 
boils  at  8°,  possesses  a  suffocating  odor,  and  is  very  soluble  in  benzene 
and  some  other  hydrocarbons.  When  brought  into  contact  with 
water  it  is  hydrolyzed  at  once,  forming  carbonic  acid  and  hydro- 
chloric acid: 

COC12  +  2H2O  ->  H2CO3  +  2HC1. 

Urea.  —  When  ammonia  and  carbonyl  chloride  are  mixed  in  the 
proper  proportions,  in  solution  in  toluene,  urea,  a  most  interesting 
chemical  substance,  is  produced: 

.Cl        H-NH2  .NH2 

0  =  C'       +  -+0  =  C(         +  2HC1 

Cl        H-NH2  NH2 

Excess  of  ammonia  has  to  be  used  to  combine  with  the  hydrogen 
chloride  thus  set  free,  so  that  the  final  equation  is : 

COC12  +  4NH3  -»  CO(NH2)2  +  2NH4C1. 

The  urea,  a  white,  crystalline  solid,  is  soluble  in  alcohol,  while  ammo- 
nium chloride  is  not,  so  that  the  former  may  be  washed  out  by  means 
of  this  solvent  and  recovered  by  evaporation.  A  little  reflection  will 
show  that,  using  the  above  action  as  the  final  stage,  urea  can  be  built 
up  from  the  simple  substances  composing  it. 

Urea  was  known  long  before  any  method  for  its  synthesis  had 
been  discovered.  It  is  the  chief  product  of  the  decomposition  of  com- 
pounds of  nitrogen  in  the  animal  body,  and  is  found  in  the  liquid 
excrements  of  animals.  It  was  regarded  as  a  typical  organic  sub- 
stance, in  the  old  sense  of  the  word  (p.  566).  In  1828,  Wohler 
succeeded  in  preparing  it  artificially  (see  below).  This  was  the  first 
synthesis  by  a  chemist  of  a  true  " organic"  substance,  and  its  prepara- 


CARBON  AND  THE  OXIDES  OF  CARBON  583 

tion  proved  to  be  the  precursor  of  many  discoveries  of  a  similar  nature. 
From  a  later  year,  about  1840,  we  may  date  the  transition  of  organic 
chemistry,  a  science  in  which  the  mystery  of  life  was  supposed  to  be 
supreme,  into  the  chemistry  of  the  compounds  of  carbon,  which  is  a 
branch  of  inorganic  chemistry. 

Wohler  used  ammonium  cyanate  (q.v.\  a  substance  in  whose  preparation  we 
are  independent  of  all  products  of  life  processes.  When  ammonium  cyanate,  or  a 
mixture  of  any  ammonium  salt  with  potassium  cyanate  in  solution  in  water,  is 
warmed  for  some  time,  an  intramolecular  change  (cf.  p.  20)  takes  place,  and  long 
prisms  of  urea  are  deposited  as  the  liquid  cools: 

NH4.CNO^±CO(NH2)2. 

Since  the  action  is  reversible,  about  four  or  five  per  cent  of  the  ammonium  cyanate 
remains  unchanged. 

The  two  substances  just  mentioned  are  entirely  different  in  chemical  proper- 
ties. Ammonium  cyanate  is  a  highly  ionized  salt,  while  urea  is  not  a  salt  at  all, 
but  a  substance  like  ammonia  which  unites  with  acids  to  form  salts.  Materials 
which,  like  these,  have  the  same  composition  and  the  same  numbers  of  units  in 
their  molecules,  and  yet  possess  different  properties,  are  spoken  of  in  chemistry  as 
isomers.  The  formulae  we  have  employed  attempt  to  explain  the  differences  in 
their  properties  by  suggesting  a  difference  in  their  molecular  structure  (cf.  p.  442). 

Assisted  by  the  catalytic  action  of  certain  ferments,  urea,  when 
dissolved  in  water,  can  take  up  two  molecules  of  the  solvent  to  form 
ammonium  carbonate: 

CO(NH2)2  +  2H2O  -»  (NH4)2C03  <=±  2NH3  +  H20  +  C02. 

Ammonium  carbonate  (q.v.)  is  an  unstable  compound,  and  in  turn, 
gives  off  ammonia  and  carbon  dioxide.  To  this  action  is  due  in 
part  the  strong  odor  of  ammonia  arising  from  the  decomposition  of 
sewage. 

Exercises.  —  1.  To  which  of  the  factors  in  the  interaction  of 
calcium  carbonate  and  hydrochloric  acid  (p.  573)  is  due  the  forward 
displacement  of  all  the  equilibria? 

2.  What  will  be  the  excess  of  pressure  inside  a  bottle  of  soda- 
water  when  four  volumes  of  carbon  dioxide  are  dissolved  in  one 
volume  of  water? 

3.  What  volume  of  liquid  carbon  dioxide,  measured  at  0°,  will  be 
required  to  give  75  liters  of  the  gas  at  0°  and  760  mm.  pressure? 

4.  What  will  be  the  effect  of  increase  in  pressure  on  the  dissocia- 
tion of  carbon  dioxide  (p.  574)? 


584  INORGANIC  CHEMISTRY 

5.  Prepare  a  diagram  showing  the  whole  scheme  of  equilibria  in- 
volved in  the  hydrolysis  of  sodium  carbonate  (p.  576). 

6.  What  volume  of  carbon  dioxide  at  0°  and  760  mm.  is  required 
for  complete  interaction  with  one  liter  of  normal  sodium  hydroxide? 

7.  What  are  the  exact  relative  weights  of  equal  volumes  of 
carbon  dioxide,  carbon  monoxide,  air,  and  steam? 


CHAPTER   XXIX 
THE  HYDROCARBONS.     ILLUMINANTS.     FLAME 

THE  compounds  of  carbon  and  hydrogen  are  called  the  hydro- 
carbons. Petroleum  is  a  mixture  of  many  substances  of  this  class. 
The  hydrocarbons  are,  therefore,  of  great  importance  in  connection 
with  fuel,  illumination,  and  lubrication. 

THE  HYDROCARBONS 

More  than  two  hundred  and  fifty  compounds  of  carbon  and  hydro- 
gen have  been  described.  They  fall  into  several  distinct  series,  the 
chief  one  of  which  contains  methane  CH4  as  its  simplest  member. 
On  account  of  the  fact  that  certain  members  of  this  set  are  found  in 
paraffin,  it  is  commonly  known  as  the  paraffin  series.  For  the  reason 
that  in  this  series  the  carbon  has  all  its  four  valences  employed,  the 
members  are  also  called  the  saturated  hydrocarbons. 

Paraffin  or  Saturated  Series  of  Hydrocarbons.  —  The  fol- 
lowing list  gives  the  formulae  of  a  few  members  of  this  series,  with 
their  names,  and  the  boiling-points  of  seven  of  the  simplest  hydro- 
carbons, and  of  two  of  the  higher  members  of  this  series  : 

Methane  CH4  b.-p.  —164°  Hexane  C6Hi4  b.-p.     71° 

Ethane      C2He     "     -   89.5°  Heptane  C7H16      "       99° 

Propane    CaHs     "     -   37°  Hexadecane          C16H34     "  '  287.5° 
Butane      C4H10    "      +     1°  "  "     m.-p.     18° 

Pentane    C6Hi2    "  35°  Pentatricontane  CasH™     "       74.7° 


After  the  first  four,  the  names  are  based  on  the  Greek  numerals 
representing  the  number  of  carbon  atoms  in  the  molecule.  Hep- 
tane is  followed  by  octane  CgHis,  nonane  CgH2o,  decane  CioH^, 
etc.  On  examining  the  formulae,  we  perceive  that,  in  each,  the 
number  of  hydrogen  atoms  is  equal  to  twice  the  number  of  carbon 
atoms  plus  two.  The  general  formula  is  therefore  CnH2n+2.  Sub- 
stances related  in  this  way  form  an  homologous  series.  The  series 
illustrates  strikingly  the  law  of  combining  weights  (p.  61).  We 

585 


586 


INORGANIC  CHEMISTRY 


note,  also,  that  the  first  four  are  gases  at  room  temperature.  The 
members  from  pentane  to  pentadecane  Ci5H32  are  liquids,  and  from 
hexadecane  onward  they  are  solids. 

In  these  compounds  the  carbon  is  quadrivalent,  and  each  sub- 
stance is  related  to  the  preceding  one  by  containing  the  additional 
unit  CH2.  The  graphic  formulae  of  the  first  three  members  illus- 
trate these  two  facts: 


H 
I 

H-C-H 
I 
H 


H      H 

I  I 
H-C-C-H 

I  I 
H  H 


H      H      H 
I        I        I 

H-C-C-C-H 

I        I        I 

H      H     H 


Transferences  of  H  one  step  to  the  right  and  interpositions  of  CH2 
constitute  the  successive  differences. 

Petroleum.  —  Petroleum  is  a  thick,  often  greenish-brown  col- 
ored oil.  When  borings  reach  the  oil-bearing  strata,  the  oil,  hitherto 
held  beneath  impervious  strata,  and  often  under  hydrostatic  pressure 
of  water  underneath  or  around  it,  either  gushes  up  or  is  pumped  to 
the  surface.  Wells  are  in  operation  in  Caucasia,  Galicia,  India, 
Japan,  and  in  Ontario,  Ohio,  Pennsylvania,  California,  Oklahoma, 
and  elsewhere.  The  world's  production  in  1912  was  400  million 
barrels  (42  gal.  each),  of  which  266  millions  were  produced  in  the 
United  States.  In  oil-refining,  advantage  is  taken  of  the  differences 
in  the  boiling-points  to  make  a  partial  separation  of  the  components 
by  fractional  distillation  (see  below) .  The  compounds  containing  sul- 
phur which  are  often  present,  and  would  give  the  obnoxious  sulphur 
dioxide  when  the  oil  was  burned,  are  deprived  of  this  constituent 
by  heating  the  oil  with  powdered  cupric  oxide  (Frasch  process). 
The  unsaturated  hydrocarbons  (q.v.)  are  removed  by  agitation  with 
concentrated  sulphuric  acid.  The  following  are  some  of  the  products 
of  the  oil  refinery,  with  their  components  and  uses. 


Name. 

Components. 

B.-P. 

Uses. 

Petroleum  ether    . 
Gasoline  (petrol)  . 
Naphtha  
Benzine    
Kerosene     .... 

Peritane-hexane 
Hexane-heptane 
Heptane-octane 
Octane-nonane 
Decane-hexadecane 

40°-  70° 
70°-  90° 
80°-120° 
120°-150° 
150°-300° 

Solvent,  gas-making 

"        fuel 

<(           a 

(t           tf 
Illuminating-oil 

THE  HYDROCARBONS.      ILLUMINANTS.      FLAME          587 

The  portions  of  still  higher  boiling-point  are  used  as  lubricating 
oils,  and  the  residue  for  fuel.  Special  treatment,  such  as  super- 
heating the  vapor  under  high  pressure  (Rittman's  process),  is  used 
to  increase  the  proportion  of  gasoline  (petrol)  for  which  there  is  a 
large  and  increasing  demand. 

The  vapor  of  these  products  is  more  inflammable  the  more  volatile 
the  components.  The  sale  of  kerosene  is  controlled  legally  by  the  re- 
quirement that  the  vapor  it  gives  when  heated  shall  not  catch  fire 
from  a  naked  flame  until  the  oil  has  reached  a  certain  minimum  tem- 
perature, the  " flash-point."  This  varies  from  37.7°  to  68.5°  in 
different  "states  and  countries.  To  be  used  safely,  the  flash-point 
should  be  65°  (150°  F). 

At  some  suitable  stage,  the  residual  oil  is  chilled,  and  a  quantity 
of  the  solid  members  of  the  series  (C22H46  to  C28H58)  crystallizes  in 
flakes  (solid  paraffin)  and  is  separated  by  filtration  in  presses.  The 
paraffin  is  used  in  waterproofing  paper,  in  laundry  work,  and  as  an 
ingredient  in  candles.  In  some  cases  petrolatum  (vaseline),  consist- 
ing of  substances  melting  at  40°-50°,  C^EUc  to  C^H^,  is  obtained 
also. 

From  ozocerite,  which  is  a  sort  of  natural  paraffin,  ceresin,  a  sub- 
stitute for  beeswax,  is  made.  Asphalt  is  another  natural  mixture  of 
solid  hydrocarbons,  found  particularly  in  Trinidad,  and  used  in  road- 
making. 

Thje  formation  of  these  hydrocarbons  in  nature  is  not  yet  thor- 
oughly explained.  According  to  one  theory,  they  are  formed  by  the 
action  of  water  upon  carbides  of  metals;  while  according  to  another, 
they  result  from  the  decomposition  of  vegetable  or  animal  matter. 
Possibly  both  of  these  sources  have  contributed  to  their  formation. 
Certain  differences  between  the  natural  oils  of  different  localities, 
such  as  the  presence  of  aromatic  hydrocarbons  in  California,  point, 
at  all  events,  to  some  difference  in  their  origin  or  subsequent  treat- 
ment. 

Fractional  Distillation.  —  When  the  boiling-points  of  two  com- 
ponents of  a  liquid  are  very  far  apart,  the  vapor  pressure  of  the  one 
may  be  very  low,  when  that  of  the  other,  by  heating,  has  reached  760 
mm.  In  this  case  the  first  distillate  will  contain  little  of  the  high- 
boiling  component.  When,  as  in  the  case  of  petroleum,  the  differ- 
ences in  boiling-points  are  not  great,  complete  separation  of  the 
components  is  difficult.  Yet  by  distillation  in  which  the  distillate 
is  caught,  not  in  one  vessel,  but  in  several  successively,  " fractions" 


588  INORGANIC   CHEMISTRY 

are  obtained  such  that  the  earlier  ones  contain  more  of  the  low-boiling 
and  the  later  ones  more  of  the  high-boiling  materials.  The  distillate 
is  diverted  into  different  vessels  when  the  thermometer  immersed  in 
the  vapor  (Fig.  20,  p.  43)  reaches  certain  temperatures,  or  (in  petroleum 
refineries)  when  the  density  of  the  distillate  reaches  certain  values, 
so  that  fractions  of  the  same  kind  are  kept  together.  When  these 
fractions  are  then  distilled  one  at  a  time,  beginning  with  the  lowest, 
and  the  several  distillates  are  divided  from  one  another  by  the  same 
temperatures  as  before,  a  more  complete  separation  is  effected. 
This  process  is  called  fractional  distillation,  and  may  be  repeated  as 
often  as  we  please  with  constantly  increasing  differentiation  of  the 
fractions. 

An  experimental  illustration  may  be  given  by  mixing  0.4  c.c.  of 
benzene  (b.-p.  80.4°)  with  8  c.c.  formic  acid  (b.-p.  100°)  and  2  c.c. 
benzyl  alcohol  (b.-p.  206.5°),  and  boiling  a  part  of  the  mixture  in  a 
test-tube  with  a  small  flame.  The  components  come  off  in  succession, 
and  are  recognized  by  the  fact  that  the  first  and  last  burn  with  a 
luminous  flame,  while  the  flame  of  the  second  is  non-luminous 
[Lect.  exp.].  By  passing  the  vapors  into  a  condenser,  and  using  the 
method  described  above,  a  more  or  less  complete  separation  can  be 
made. 

General  Properties  of  Paraffins.  —  All  these  substances  are 
extremely  indifferent  in  their  chemical  behavior.  They  have  none 
of  the  properties  of  acids,  bases,  or  salts.  The  halogens,  notably 
chlorine  and  bromine,  however,  interact  with  them  (see  below). 
When  burned  they  all  produce  carbon  dioxide  and  water.  When  their 
vapors  are  passed  through  a  white-hot  tube  they  suffer  decomposition 
into  a  mixture  of  hydrogen  and  hydrocarbons  of  smaller  or  larger 
(see  Benzene)  molecular  weight. 

Methane  CH^.  —  Methane,  otherwise  known  as  marsh-gas,  is 
the  chief  component  of  natural  gas.  It  rises  to  the  surface  when  the 
bottoms  of  marshy  pools  are  disturbed,  and  issues  from  seams  in  coal 
beds.  In  these  two  cases  it  results  from  the  decomposition  of  vege- 
table matter  in  absence  of  air.  When  methane  enters  mines  from 
a  coal  seam  it  is  called  "fire-damp"  (Ger.  Dampf,  vapor),  on  ac- 
count of  the  explosive  nature  of  the  mixture  it  forms  with  the  air. 
The  carbon  dioxide  formed  by  the  explosion  is  called  by  the  miners 
"  choke-damp. "  In  many  regions  it  is  confined,  like  the  oil,  beneath 
impervious  strata  and  is  forced  out  through  borings  by  hydro- 


THE  HYDROCARBONS.      ILLUMINANTS.      FLAME          589 

static  pressure.     It  is  found  mainly  in  or  near  the  localities  where 
oil  is  found. 

The  formation  of  methane  by  direct  union  of  carbon  and  hydro- 
gen has  already  been  discussed  (p.  570).  It  may  be  made  from 
inorganic  materials  by  the  action  of  water  upon  aluminium  carbide, 
prepared  by  the  interaction  of  aluminium  oxide  and  carbon  in  the 
electric  furnace  (cf.  p.  569)  : 

A14C3  +  12H20  -»  4A1(OH)3  +  3CH4  |. 

In  the  laboratory  the  gas  is  commonly  obtained  by  the  distillation 
of  a  dry  mixture  of  sodium  acetate  and  sodium  hydroxide  : 

NaC02CH3  +  NaOH  ->  Na^COs  +  CH4  T. 


As  regards  chemical  properties,  methane,  like  other  saturated 
hydrocarbons  (p.  588),  is  very  indifferent.  When  a  mixture  of 
methane  and  chlorine  is  exposed  to  sunlight  several  changes  occur  in 
succession  (cf.  p.  225): 

CH4  +  C12  ->  CH3C1  +  HC1,        CH2C12  +  C12  -»  CHC13  +  HC1, 
CH3C1  +  C12  -*  CH2C12  +  HC1,       CHC13  +  C12  -*  CC14  +  HC1. 

This  kind  of  interaction  with  the  halogens  is  characteristic  of 
the  saturated  hydrocarbons.  It  takes  place  slowly,  and  is  there- 
fore entirely  different  from  ionic  chemical  change.  It  consists  in  a 
progressive  substitution  of  chlorine  for  hydrogen,  unit  by  unit.  The 
various  groups  which,  in  the  first  three  of  these  products,  are  asso- 
ciated with  chlorine,  occur  in  many  organic  compounds,  and  receive 
the  names  methyl  CH3—  ,  methylene  CH2  =  ,  and  methenyl  CH=. 
The  compounds  are  known,  therefore,  as  methyl  chloride,  methylene 
chloride,  methenyl  chloride  (chloroform),  and  carbon  tetrachloride 
(p.  570).  The  last  two  are  volatile  liquids,  and  familiar  substances. 
The  corresponding  iodine  derivative  iodoform  CHI3  is  used  in  surgical 
dressing.  These  substances  are  not  salts,  and  are  not  ionized  in 
solution.  They  are  very  slowly  hydrolyzed  by  water  —  carbon 
tetrachloride,  for  example,  giving  carbonic  acid  and  hydrochloric 
acid.  Although  carbon  is  a  non-metal  (cf.  p.  209),  this  action  requires 
a  high  temperature.  Methane  and  the  other  saturated  hydrocarbons 
are  decomposed  by  strong  heating  (see  cracking,  below). 

Organic  Radicals.  —  In  carbon  chemistry  there  are  groups  of 
units  which  pass  unaltered  from  compound  to  compound  and  receive 
the  name  organic  radicals.  They  usually  lack  a  property  which  inor- 


590  INORGANIC  CHEMISTRY 

ganic  radicals  generally  possess,  namely,  the  power  to  form  ions 
(p.  324).  Methyl  is  such  a  radical,  being  found  in  methane  CH3.H, 
methyl  chloride  CH3.C1,  methyl  alcohol  CH3.OH,  and  acetic  acid 
CH3.CO2H.  Similarly  we  have  ethyl  CzH-5  in  ethane  C2H5.H  and  in 
ethyl  alcohol  C2H5OH.  Methyl,  ethyl,  and  propyl  C3H7—  are  uni- 
valent  radicals.  We  have  also  ethylene  C2H4  = ,  propylene  CsHe  = , 
and  so  forth,  which  are  bivalent.  Groups  like  NO2J  (p.  527),  NH2X 
(p.  519),  CHsCO1,  and  many  more,  are  other  non-ionizable  radicals 
found  in  organic  compounds  (see  Acetic  acid,  below). 

Unsaturated  Hydrocarbons.  —  In  addition  to  the  saturated 
series  of  hydrocarbons,  several  other  series  are  known  in  which 
smaller  proportions  of  hydrogen  are  present.  Thus,  ethylene  C2H4, 
to  which  illuminating  gas  largely  owes  the  luminosity  of  its  flame, 
belongs  to  a  series  CnH2n,  all  the  members  of  which  contain  two 
atoms  of  hydrogen  less  than  the  corresponding  compounds  of  the 
first  series.  Again,  acetylene  C2H2  is  the  first  member  of  a  series 
CnH2n_2,  and  benzene  C6H6  begins  a  series  CnH2n_6,  of  which  toluene 
CyHs  (p.  527)  is  the  second  member.*  These  are  all  unsaturated 
because  the  full  valence  of  the  carbon  is  not  in  use,  and  these  com- 
pounds, therefore,  unite  more  or  less  readily  with  hydrogen,  chlorine, 
bromine,  and  concentrated  sulphuric  acid.  The  hydrocarbons 
of  all  the  series  are  mutually  soluble,  but  none  of  them  dissolves  in 
water. 

Members  of  the  ethylene  and  acetylene  series  are  found  in  petro- 
leum, and  are  formed  also  to  some  extent  by  decomposition  during 
the  distillation.  As  oil  containing  them  acquires  dark-colored 
products  by  chemical  change,  the  oils  are  always  refined  before 
being  sold.  They  are  agitated  with  concentrated  sulphuric  acid, 
which  unites  with  the  unsaturated  substances  and,  being  insoluble 
in  the  oil,  collects  in  a  layer  below  it.  The  oil  is  finally  washed 
free  from  the  acid  with  dilute  alkali  and  with  water. 

Ethylene  C2H4. —  Ethylene  is  the  first  member  of  the  second 
series  of  hydrocarbons.  It  corresponds  to  ethane  C2H6,  but  con- 
tains in  each  molecule  two  hydrogen  units  less  than  does  this  sub- 
stance. 

*  Isoprene  CsHs,  a  member  of  the  unsaturated  series  CnH2n-4,  when  heated 
in  presence  of  sodium  (or  some  other  contact  agent),  changes  into  caoutchouc 
(CsHs)*  or  raw  rubber.  No  method  of  preparing  artificial  rubber  has  yet  been 
used  commercially. 


THE  HYDROCARBONS.      ILLUMINANTS.      FLAME          591 

Ethylene  is  made  by  heating  common  alcohol  (ethyl  alcohol) 
with  concentrated  sulphuric  acid: 

C2H5OH-»H20  +  C2H4T. 

The  action  with  sulphuric  acid  really  takes  place  in  two  distinct  stages, 
and  the  intermediate  product  can  be  isolated.  First,  ethyl  hydro- 
gen sulphate  (cf.  p.  385)  is  formed,  C2H5OH  +  H2SO4  ^  C2H5HS04 
+  H2O.  Above  150°,  however,  this  substance,  which  is  a  thick 
syrup,  is  dissociated,  giving  ethylene  and  sulphuric  acid,  C-jHsHSC^ 
— -»  C2H4  +  H2S04.  A  comparison  of  the  structural  formulae  of  the 
alcohol  and  ethylene  shows  that  this  loss  of  water  must  leave  the 
carbon  partly  unsaturated: 

H     H  H     H  H    H 

II  II  II 

H-C-C-O-H  H-C-C-H  or    H-C  =  C-H 

II  II 

H     H 

The  water  may  also  be  removed  by  allowing  alcohol  to  fall  drop  by 
drop  on  heated  phosphoric  anhydride.  The  solid  metaphosphoric 
acid  remains  behind  and  ethylene  escapes. 

Ethylene  is  formed  along  with  acetylene  and  other  substances, 
when  any  saturated  hydrocarbon  is  heated  strongly.  Even  methane 
gives  it: 

2CH4  ->  C2H4  +  2H2. 

Ethylene  is  a  gas,  which,  when  liquefied,  boils  at  —105°.  Its 
critical  temperature  is  35°.  At  0°  it  may  be  liquefied  by  a  pressure  of 
42  atmospheres.  It  burns  in  the  air  with  a  flame  which,  on  account  of 
the  great  separation  of  free  carbon  which  takes  place  temporarily  dur- 
ing the  combustion  (cf.  Flame),  is  highly  luminous.  It  will  be  seen 
that  in  the  formula  but  three  of  the  valences  of  each  carbon  unit  are 
occupied.  As  carbon  is  usually  either  bivalent  or  quadrivalent,  we 
should  expect  that  in  this  compound  the  combining  capacity  of  the 
carbon  would  not  be  completely  satisfied.  We  find  this  to  be  the 
case.  Ethylene  is  easily  reduced  by  active  hydrogen  (p.  543)  to 
ethane,  taking  up  two  units  of  hydrogen  in  the  process.  When 
ethylene  is  passed  through  liquid  bromine,  it  is  rapidly  absorbed,  and 
the  bromine  seems  to  increase  in  volume  and  finally  loses  all  its  color, 
leaving  a  transparent  liquid  having  the  composition  C2H4Br2,  ethyl- 
ene bromide.  The  second  of  the  above  graphic  formulae  for  ethylene 


592  INORGANIC  CHEMISTRY 

is  the  one  generally  used.  In  spite  of  appearances,  it  is  not  in- 
tended to  indicate  that  the  two  units  of  carbon  are  more  forcibly  held 
together  than  in  other  compounds  (cf.  p.  138).  It  simply  chronicles 
the  fact  that  one  valence  of  each  carbon  unit  is  unoccupied. 

Acetylene.  —  This  substance,  likewise  a  gas,  is  the  first  member 
of  still  another  unsaturated,  homologous  series,  CnH27l_2.  Its  for- 
mula C2H2  shows  that  its  molecule  lacks  four  of  the  hydrogen  units 
necessary  to  the  complete  saturation  which  we  find  in  ethane. 
Graphically  its  structure  is  usually  represented  thus:  H  —  C  =  C  —  H. 
This  gas  is  formed  in  small  quantities  by  direct  union  of  carbon  and 
hydrogen  in  the  electric  arc  (p.  570).  It  is  formed,  not  because  the 
elements  prefer  to  give  this  one  of  the  many  hydrocarbons,  but  be- 
cause, at  2000°  or  over,  the  other  hydrocarbons  are  decomposed,  and 
acetylene  is  one  in  the  formation  of  which  the  greatest  amount  of 
heat  is  absorbed  (van't  Hoff's  law,  p.  305,  and  see  below).  For  the 
same  reason,  it  is  produced  when  ethylene  is  passed  through  a  heated 
tube:  C2H4-*  C2H2  +  H2  (cf.  Flame). 

When  calcium  carbide  (p.  571)  is  thrown  into  water,  violent 
effervescence  occurs,  the  carbide  is  disintegrated,  a  precipitate  of 
calcium  hydroxide  is  formed,  and  acetylene  passes  off  as  a  gas : 

CaC2  +  2H2O  ->  Ca(OH)2  +  C2H2  T- 

This  hydrolytic  action  is  like  that  of  water  on  calcium  phosphide 
(p.  553),  calcium  sulphide  (p.  421),  and  magnesium  nitride  (p.  514). 

Acetylene  burns  with  a  flame  which  is  still  more  luminous  than 
that  of  ethylene.  Its  most  characteristic  property  is  that  when  passed 
through  an  ammoniacal  solution  of  a  cuprous  salt,  it  yields  a  red  pre- 
cipitate of  a  carbide  of  copper  known  as  copper  acetylide.  The 
equation:  Cu2(OH)2  +  C^E2  — >  Cu2C2  +  2H2O,  partially  represents 
the  change.  This  red  precipitate,  when  dried,  is  extremely  explosive, 
on  account  of  the  great  amount  of  energy  set  free  when  it  breaks  up 
into  its  constituents.  Its  formation  is  used  as  a  test  for  acetylene  in 
mixtures  of  gases. 

Acetylene  may  be  handled  safely  as  a  gas  at  the  ordinary  pressure, 
but,  when  contained  in  cylinders  at  more  than  two  atmospheres 
pressure,  it  is  readily  exploded  by  any  shock.  This  is  due  to  the  fact 
that  it  is  an  endothermal  compound:  C2H2  — >  2C  +  H2  -f  53,200  cal. 
It  is  frequently  made  in  generators,  as  needed,  from  calcium  car- 
bide, and  is  used  for  lighting  on  automobiles  and  in  regions  remote 
from  a  public  supply  of  illuminating-gas.  The  acetylene  tanks,  which 


THE  HYDROCARBONS.      ILLUMINANTS.      FLAME          593 

are  also  in  use,  contain  acetylene  dissolved  under  high  pressure 
in  acetone,  a  form  in  which  it  can  be  handled  safely. 

When  acetylene  C2H2  is  burned,  we  obtain  from  2  X  12  +  2  = 
26  g.  not  only  the  heat  due  to  the  combustion  of  the  carbon  (2  X 
12  X  8040  cal.,  p.  567),  and  of  the  hydrogen  (2  X  28,800  cal.),  but 
also  the  heat  due  to  the  decomposition  of  the  gas  (53,200  cal.). 
The  temperature  of  the  flame  is  therefore  extraordinarily  high. 
The  oxyacetylene  flame,  produced  by  means  of  a  suitable  burner 
(Fig.  51,  p.  126),  is  now  used,  under  the  name  of  the  acetylene 
torch,  for  cutting  metals.  The  gases  are  contained  in  portable 
tanks.  Such  a  flame  will  melt  its  way  through  a  6-inch  shaft  of 
steel  plate  several  feet  wide  in  less  than  a  minute,  cutting  the  object 
in  two.  Steel  buildings  have  thus  been  taken  down,  and  ships 
(like  the  Maine)  have  been  cut  up  for  removal.  Other  gases,  like 
blau  gas  and  oil  gas,  made  by  cracking  petroleum  (see  below),  are 
now  displacing  acetylene  for  this  purpose,  as  they  are  almost  as 
effective,  and  the  flame  is  more  easily  controlled. 

The  unsaturated  nature  of  this  substance  is  shown  by  the  avidity 
with  which  it  unites  with  hydrogen  and  the  halogens,  forming  satu- 
rated compounds. 

Benzene  CGH&.  —  Limits  of  space  forbid  the  discussion  of  any 
of  the  other  series  of  hydrocarbons.  One  of  the  most  important  has 
not  been  mentioned,  however.  It  is  that  of  which  the  first  member  is 
benzene,  CeH6.  More  than  half  of  the  known  compounds  of  carbon 
are  derived  from  this  substance.  Phenol  (cf.  p.  527)  CcHsOH  is  the 
fundamental  alcohol  of  this  set.  Benzene  is  obtained  from  the  prod- 
ucts of  the  dry  distillation  of  coal  (cf.  Goal  gas) ,  being  formed,  prob- 
ably, from  the  acetylene  which  the  decomposition  of  other  hydro- 
carbons yields.  At  all  events,  when  acetylene  is  passed  through  a 
heated  tube  some  benzene  is  produced,  3C2H2  — •>  CeHe,  along  with 
free  carbon  and  hydrogen.  Toluene  (p.  527)  CeHsCHs  is  the  second 
member  of  this  series. 

Cracking  of  Hydrocarbons.  —  All  hydrocarbons,  when  heated 
strongly  (air-excluded)  decompose  or  crack.  The  changes  seem 
to  be  reversible,  and  the  result  therefore  depends  upon  the  con- 
ditions. Thus,  at  atmospheric  pressure,  and  especially  when  the 
oil  is  mainly  present  as  a  liquid,  [hydrogen  is  given  off  and  un- 
saturated liquid  and  gaseous  hydrocarbons  are  produced.  Under 
such  conditions,  ethylene  is  formed  in  large  amounts.  On  the 


594  INORGANIC  CHEMISTRY 

other  hand,  when  an  oil  free  from  gasoline  is  completely  vaporized 
(500°),  and  is  under  high  pressure,  the  hydrogen  is  forced  into 
combination  with  the  broken  molecules  and  the  saturated  con- 
stituents of  gasoline  are  produced  in  large  amount  (Rittman's 
process). 

At  a  white  heat,  all  the  hydrocarbons  decompose  into  hydrogen 
and  free  carbon.  The  latter  is  deposited  in  a  dense  form  called 
gas-carbon,  which  is  more  or  less  crystalline  (like  graphite)  and 
used  in  making  carbon  rods  for  arc  lights  and  electric  furnaces, 
and  carbon  plates  for  batteries,  and  for  the  electrodes  employed 
in  electrolysis.  The  carbon  is  ground  up,  moistened  with  petro- 
leum residues,  subjected  to  hydraulic  pressure  and  finally  heated 
strongly  to  expel  volatile  matter. 

Carburetted  Water  Gas.  —  As  we  have  seen,  water  gas  is 
essentially  H2  +  CO  (p.  577),  and  burns  with  a  pale-blue  flame. 
To  fit  it  for  use  as  illuminating-gas,  unsaturated  hydrocarbons, 
which  burn  with  a  luminous  flame,  such  as  ethylene  C2H4  and 
acetylene  C2H2  must  be  added.  The  gas  is  sent  through  a  tower 
containing  strongly  heated  brick  on  which  a  petroleum  oil  is  sprayed. 
Mixed  with  the  vapor,  the  gas  then  passes  into  the  "superheater" 
where,  at  a  higher  temperature,  the  cracking  into  unsaturated 
hydrocarbons  occurs.  The  gas  is  then  cooled  and  washed  to  remove 
condensible  hydrocarbons,  which  would  otherwise  obstruct  the 
service  pipes.  A  typical  carburetted  water  gas  has  the  composition : 
Illuminants  17  per  cent;  heating  gases,  methane  20  per  cent,  hydrogen 
32  per  cent,  carbon  monoxide  26  per  cent;  impurities  (nitrogen  and 
carbon  dioxide)  5-6  per  cent.  A  flame  burning  5  cu.  ft.  per  hour 
gives  25  candle  power. 

Blau  gas  and  Oil  gas,  such  as  Pintsch  gas,  contain  larger  pro- 
portions of  illuminants.  Thus  a  good  gas  shows:  illuminants 
45  per  cent;  heating  gases,  methane  39  per  cent,  hydrogen  14.5 
per  cent;  impurities  1.5  per  cent;  candle  power  65.  Such  gases 
are  compressed  in  tanks  and  used  for  illumination  on  railway  trains 
(Coal  gas,  see  p.  612). 

FLAME 

Meaning  of  the  Term.  —  In  the  combustion  of  charcoal  there  is 
hardly  any  flame,  for  the  light  emanates  almost  entirely  from  the 
incandescent,  massive  solid.  When  two  gases  are  mixed  and  set  on 


THE  HYDROCARBONS.      ILLUMINANTS.      FLAME 


595 


Fia.  129. 


fire,  a  sort  of  flame  passes  through  the  mixture,  but  this  can  hardly 
be  accounted  a  flame,  in  the  ordinary  sense,  either.  The  rapid  move- 
ment of  the  flash,  and  the  explosion  which  accompanies  it,  are  in  a 
manner  the  precise  opposite  of  the  quiet  combustion 
which  is  characteristic  of  flames. 

With  illuminating-gas  the  production  of  its  very 
characteristic  flame  is  due  to  the  chemical  union  of 
a  stream  of  one  kind  of  gas  in  an  atmosphere  of  an- 
other. The  flame  is  made  up  of  the  heated  matter 
where  the  two  gases  meet.  In  the  case  of  a  burning 
candle  (Fig.  129),  *^ne  of  the  bodies  appears  to  be  a 
solid,  but  a  closer  scrutiny  of  the  phenomenon  shows 
that  the  solid  does  not  burn  directly.  A  combustible 
gas  is  manufactured  continuously  by  the  heat  of  the 
combustion  and  rises  from  the  wick.  The  introduction  of  a  narrow 
tube  into  the  interior  of  the  flame  enables  us  to  draw  off  a  stream  of 
this  gas  and  to  ignite  it  at  a  remote  point.  Thus,  a  flame  is  a  phe- 
nomenon produced  at  the  surface  where  two  gases  meet  and  undergo 
combination  with  the  evolution  of  heat  and  more  or  less  light. 

In  the  chemical  point  of  view,  it  is  a  matter  of  in- 
difference whether  the  gas  outside  the  flame  contains 
oxygen,  and  the  gas  inside  consists  of  substances 
ordinarily  known  as  combustibles,  or  whether  this 
order  is  reversed.  In  an  atmosphere  of  ordinary 
illuminating-gas,  the  flame  must  be  fed  with  air. 
This  condition  is  easily  realized  (Fig.  130).  The 
lamp-chimney  is  closed  at  the  top  until  it  has  be- 
come filled  with  illuminating-gas.  After  the  lapse  of 
a  few  minutes  this  can  be  ignited  as  it  issues  from 
the  bottom  of  the  wide,  straight  tube  which  pro- 
jects from  the  interior.  When  the  hole  in  the  cover 
of  the  lamp-chimney  is  then  opened,  the  upward 
draft  causes  the  flame  of  the  burning  gas  to  recede 
up  the  tube,  and  there  results  a  flame  fed  by  air 
and  burning  in  coal-gas.  In  an  atmosphere  of 
this  kind,  materials  playing  the  part  of  a  candle  burning  in  air 
would  have  to  be  substances  which,  under  the  influence  of 
the  heat  of  combustion,  give  off  oxygen.  Strongly  heated 
potassium  chlorate,  for  example,  appears  to  burn  continuously 
in  such  an  atmosphere  when  lowered  into  it  in  a  deflagrating 
spoon. 


Fia.  130. 


596  INORGANIC  CHEMISTRY 

Luminous  Flames:  Lampblack.  —  The  flame  of  hydrogen, 
under  ordinary  circumstances,  is  almost  invisible,  nearly  all  the 
energy  of  the  combustion  being  devoted  to  the  production  of  heat.  A 
part  of  this,  however,  may  be  transformed  into  light  by  the  suspension 
of  a  suitable  solid  body,  such  as  a  platinum  wire,  in  the  flame.  The 
holding  of  a  piece  of  quicklime  in  an  oxyhydrogen  flame  (cf.  p.  126) 
is  a  practical  illustration  of  this  method  of  securing  luminosity. 
In  general,  luminosity  may  be  produced  by  the  presence  of  some  solid 
which  is  heated  to  incandescence. 

In  the  Welsbach  lamp  the  flame  itself  is  non-luminous  and,  but 
for  the  mantle,  would  be  identical  with  the  ordinary  Bunsen  flame. 
The  mantle  which  hangs  in  the  flame,  however,  by  its  incandescence, 
furnishes  the  light.  This  mantle  is  composed  of  a  mixture  of  99  per 
cent  thorium  dioxide  ThO2  and  one  per  cent  cerium  dioxide  CeO2. 
While  many  oxides  would  give  out  a  white  light,  and  could  be  obtained 
much  more  cheaply  than  these,  they  have  not  sufficient  coherence  to 
make  their  use  practicable.  Le  Chatelier  found  that  the  tempera- 
ture of  the  flame  was  1700-1800°  C.  The  brightness  of  the  light  is 
due  to  the  fact  that  light  rays  are  plentiful,  and  few  heat  rays  are 
produced.  The  Welsbach  lamp  gives  four  times  as  much  light  as 
does  the  same  gas,  issuing  at  the  same  rate,  from  an  ordinary  burner. 
It  is  worth  noting  that  any  appreciable  variation  from  the  above 
proportions,  by  the  introduction  of  either  more  or  less  cerium  oxide, 
produces  a  marked  diminution  in  the  intensity  and  whiteness  of  the 
light  (see  Thorium). 

In  cases  of  brilliant  combustion,  as  of  magnesium  ribbon  or  phos- 
phorus, a  solid  body  is  formed  whose  incandescence  accounts  for  the 
light.  The  flame  of  ordinary  illuminating-gas  does  not  at  first  sight 
appear  to  give  evidence  of  the  presence  of  any  solid  body.  But  if  a 
cold  evaporating  dish  is  held  in  the  flame  for  a  moment,  a  thick 
deposit  of  finely  divided  carbon  (soot)  is  formed,  and  we  at  once 
realize  that  the  light  is  due  to  the  glow  of  these  particles  in  a  mass  of 
intensely  hot  gas.  Carbon  is,  indeed,  an  extremely  combustible 
substance,  and  is  eventually  entirely  consumed.  But  a  fresh  supply 
is  continually  being  generated  in  the  interior  of  the  flame,  while  the 
oxygen  with  which  it  is  to  unite  is  outside  the  flame  altogether. 
Thus  the  carbon  particles  persist  until,  drifting  with  the  spreading  gas, 
they  reach  the  periphery  of  the  flame. 

On  a  large  scale,  oil  residues  are  burned  so  that  the  flame 
strikes  a  revolving,  iron  vessel  cooled  with  water.  The  soot 
or  lampblack  is  continuously  scraped  off  as  the  vessel  turns. 


THE  HYDROCARBONS.      ILLUMINANTS.     FLAME          597 

Lampblack  is  used  in  making  printer's  ink,  India  ink,  and  black 
varnish. 

It  cannot  be  said  that  no  flames  are  luminous  unless  a  solid  body  is 
contained  in  them.  When  compressed  hydrogen  is  burned  in  an 
atmosphere  of  oxygen  under  pressure,  the  light  given  out  by  the  flame 
is  much  greater,  and,  in  general,  illuminating  power  seems  to  be 
heightened  by  increase  in  the  concentration  of  the  gas.  In  special 
cases,  also,  such  as  the  explosion  of  a  mixture  of  nitric  oxide  with 
carbon  bisulphide  vapor  [Lect.  exp.],  a  flame  which  has  intense  illu- 
minating power  is  produced,  although  the  density  of  the  gases  is 
low  and  solids  are  lacking. 

The  Bunsen  Flame:  The  Blast  Lamp.  —  In  the  burner  de- 
vised by  Robert  Bunsen,  a  jet  of  ordinary  illuminating-gas  is  pro- 
jected from  a  narrow  opening  into  a  wider  tube  (Fig. 
131).  In  this  tube  it  becomes  mixed  with  air,  forced 
by  atmospheric  pressure  through  openings  whose 
dimensions  can  be  altered  by  means  of  a  perforated 
ring.  When  the  supply  of  air  is  sufficient,  the  flame 
becomes  non-luminous.  With  a  somewhat  different 
construction,  and  the  use  of  a  bellows  to  force  a  larger 
proportion  of  air  into  the  gas,  a  still  hotter  flame  can  be 
produced.  The  instrument  in  this  case  is  known  as 
a  blast-lamp. 

The  high  temperature  of  the  blast-lamp  flame  pre- 
sents an  interesting  problem.  The  same  amounts  of 
gas  and  air  burn  to  give  the  same  amounts  of  the 
same  products,  whether  the  air  blast  is  on  or  off. 
The  same  amount  of  heat  is  produced,  and  the  same 
quantities  of  the  same  substances  are  heated.  The 
average  temperature  throughout  the  flame  should 
therefore  be  the  same.  In  point  of  fact,  it  is  the  same,  but  the 
stream  of  hot  gas  is  moving  more  rapidly  when  the  blast  is  going. 
The  temperature  of  a  body  immersed  in  the  flame  depends,  on 
the  one  hand  upon  the  rate  at  which  heat  reaches  it,  and  upon 
the  other  on  the  rate  at  which  it  loses  heat  by  radiation.  The 
heat  is  partly  carried  by  the  moving,  heated  gases  (convection), 
and  partly  transmitted  by  conduction  through  the  stationary  layer 
(p.  503)  on  the  surface  of  the  body.  Now,  the  latter  is  the  slower 
process.  Hence  a  rapid  stream  of  gas,  which  leaves  a  thinner  station- 
ary layer,  will  diminish  the  distance  the  heat  has  to  travel  by  con- 


598  INORGANIC  CHEMISTRY 

duction  and  so  convey  heat  to  the  body  faster  than  could  a  slow 
stream  of  the  same  temperature.  Thus,  with  a  blast  flame,  the  loss 
by  radiation  is  the  same  at  the  same  temperature,  but  heat  reaches 
the  body  faster  and  so  the  temperature  of  the  body  more  nearly 
approaches  that  of  the  flame  itself. 

The  Bunsen  flame  is  hotter  than  an  ordinary  flame,  because  in  it, 
also,  the  gases  move  faster.  It  is  instructive  to  note  the  effect  of 
forcing  in  larger  and  larger  proportions  of  air  into  the  Bunsen  flame. 
The  flame  at  the  top  of  the  tube  continually  diminishes  in  size,  even 
after  it  has  become  non-luminous.  Finally,  a  point  is  reached  at 
which  the  flame  is  so  unstable  that  the  smallest  further  increase  in 
the  supply  of  air  causes  it  to  descend  into  the  tube.  The  mixture 
of  illuminating-gas  and  air  in  the  tube  of  a  Bunsen  burner  is  an 
explosive  one,  and  the  explosion-flame  will  proceed  through  it  with 
greater  rapidity,  the  more  nearly  the  quantity  of  air  approaches  that 
required  for  complete  combustion.  When  the  speed  with  which  the 
explosion-flame  would  move,  equals  that  with  which  the 
stream  of  the  mixed  gases  is  coming  upward  through  the 
tube,  the  flame  reaches  the  unstable  condition  just  men- 
tioned. Any  increase  in  the  proportion  of  air  raises 
the  speed  with  which  the  explosion  can  travel,  and  the 
flame  is  thus  able  to  proceed  down  the  tube  against  the 
stream  of  gas.  This  phenomenon  is  frequently  noticed 
in  the  Bunsen  burner,  when  the  holes  admitting  the  air 
are  too  large,  or  a  draft  momentarily  causes  an  increase 
in  the  supply  of  air.  The  flame  strikes  back,  and  there- 
after continues  to  burn  at  the  bottom  of  the  tube. 

Structure  of  the  Bunsen  Flame.  —  When  an  ex- 
ceedingly small  luminous  flame  is  examined,  the  various 
parts  of  which  it  consists  may  easily  be  made  out.  In 
the  interior  there  is  a  dark  cone  which  is  composed  of 
illuminating-gas  and  air,  and  in  it  no  combustion  is 
FIG  132  taking  place.  A  match-head  may  be  held  here  for  some 
time  without  being  set  on  fire.  This  is  therefore  not 
properly  a  part  of  the  flame.  Outside  this  is  a  vivid  blue  layer  (C, 
Fig.  132)  which  is  best  seen  in  the  lower  part  of  the  flame,  but  extends 
beneath  the  luminous  sheath,  and  covers  the  dark  inner  cone  com- 
pletely. Outside  the  blue  flame,  and  covering  the  greater  part  of  it, 
is  the  cone-shaped  luminous  portion  (B).  Over  all  is  an  invisible 
mantle  of  non-luminous  flame  (A),  which  becomes  visible  when  the 


THE  HYDROCARBONS.      ILLUMINANTS.      FLAME          599 

light  from  the  luminous  part  is  purposely  obstructed.  In  the  lumi- 
nous gas-flame,  therefore,  there  are  four  regions,  if  we  count  the  inner 
cone  of  gas.  The  difference  between  this  and  the  non-luminous 
Bunsen  flame  is  that  in  the  latter  the  luminous  region  is  omitted, 
and  the  inner,  dark  cone,  the  blue  sheath,  and  the  outer  mantle, 
are  the  only  parts  which  can  be  distinguished.  We  shall  see  that 
in  these  different  regions  the  chemical  changes  taking  place  are 
different. 

The   Causes   of  Luminosity  and  Non-Luminosity.  —  The 

study  of  the  chemical  changes  taking  place  in  the  Bunsen  flame, 
particularly  with  the  object  of  explaining  (1)  the  luminosity  of  the 
flame  of  the  pure  gas,  and  (2)  the  non-luminosity  of  that  produced 
by  the  same  gas  when  it  is  mixed  with  air,  has  been  the  subject  of 
many  elaborate  investigations.  The  questions  are:  (1)  Why  is 
carbon  liberated  in  the  former  case,  and  (2)  why  is  it  not  liberated  in 
the  latter?  Let  us  consider  these  questions  in  order. 

1.  The  investigations  of  Lewes  (1892)  and  others  show  conclu- 
sively that  the  free  carbon  in  the  luminous  zone  of  the  ordinary 
flame  is  accompanied  by  free  hydrogen,  and  that  both  are  formed 
by  dissociation  of  the  ethylene  in  the  inner  blue  cone.  This  sub- 
stance, when  heated,  gives  acetylene,  and  the  latter  then  dissociates 
into  carbon  and  hydrogen  (p.  591) : 

O2H.4  — ^  tl2     I     02X12  — ^  2O  ~j~  Jl2. 

The  carbon  glows,  until,  as  it  drifts  outwards,  it  encounters  the 
oxygen  of  the  air  and  is  burned.  The  first  oxygen  encountered 
combines  more  readily  with  the  hydrogen,  since  it  is  a  gas,  than 
with  the  carbon,  which  is  now  in  solid  particles  and  therefore  burns 
less  readily.  That  carbon  glows  when  heated  in  the  absence  of 
oxygen,  without  being  consumed,  is  a  fact  familiar  in  the  behavior 
of  the  incandescent  electric  lamp,  the  filaments  of  which  were  all 
formerly  made  of  carbon. 

The  conception  that  when  hydrocarbons  burn,  they  first  undergo 
dissociation,  and  then  union  with  oxygen,  is  in  harmony  with  what 
we  have  observed  also  in  the  case  of  the  combustion  of  hydrogen 
sulphide,  where  the  presence  of  free  sulphur  and  free  hydrogen  in 
the  interior  of  the  flame  was  demonstrated  (p.  416).  That  acetylene 
is  actually  formed  as  an  intermediate  substance  is  easily  shown. 
It  is  found  that  when  the  Bunsen  flame  " strikes  back,"  and  the 
combustion  of  the  gases  is  rendered  incomplete  by  the  contact  of  the 


600 


INORGANIC  CHEMISTRY 


flame  with  the  cold  tube,  a  large  amount  of  acetylene  is  formed. 
Again,  when  the  gases  surrounding  the  flame  of  air  burning  in 
illuminating-gas  (p.  595)  are  withdrawn  by  means  of  a  pump  and 
caused  to  pass  through  an  ammoniacal  solution  of  cuprous  chloride 
(Fig.  133),  large  quantities  of  copper-acetylide  are  precipitated. 

2.    The  influence  of  the  air  admitted  to  the  Bunsen  burner,  in 
interfering  with  this  dissociation  in  such  a  way  as  to  destroy  all 


FIG.  133. 

luminosity,  is  the  most  difficult  point  to  explain.  The  effect  is 
frequently  attributed  to  the  oxygen  which  the  air  contains.  This 
view,  however,  is  seriously  weakened  by  a  consideration  of  the  un- 
doubted fact  that  oxygen  is  not  required.  Carbon  dioxide  or 
steam  is  equally  efficient  when  introduced  instead  of  air  (Fig.  134, 
gas  enters  at  a  and  C02  at  6).  Even  nitrogen,  which  cannot  possibly 
be  suspected  of  furnishing  any  oxygen,  likewise  destroys  the  lumi- 
nosity. Lewes  has  shown  that  0.5  volumes  of  oxygen  in  1  volume  of 
coal  gas  destroy  the  luminosity.  But  2.30  volumes  of  nitrogen  or 
2.27  volumes  of  air  accomplish  the  same  result.  Thus  the  efficiency 
of  air  is  not  much  greater  than  that  of  nitrogen,  in  spite  of  the  fact 
that  one-fifth  of  the  former  is  oxygen. 

It  is  evident  that  the  effect  is  due,  in  part  at  least,  to  the  dilution 
with  a  cold  gas.  This  is  confirmed  by  the  observation  that  a  cold 
platinum  dish  held  in  a  small  luminous  flame  is  similarly  destruc- 
tive of  the  luminosity.  If  the  tube  of  the  Bunsen  burner  is  heated, 
so  that  the  mixed  gases  are  considerably  raised  in  temperature 
before  reaching  the  non-luminous  flame,  the  latter  becomes  In  mi- 


THE  HYDROCARBONS.      ILLUMINANTS.     FLAME 


601 


nous.  It  is  probable,  therefore,  that  the  cold  gas  lowers  the  tem- 
perature of  the  inner  flame,  and  at  the  same  time  the  dilution 
diminishes  the  speed  with  which  the  free  carbon  is  formed  (Lewes). 
Even  if  the  temperature  is  not  reduced  below  that  at  which  dis- 
sociation of  the  ethylene  can  occur,  yet  the  dilution  and  cooling, 
together,  prevent  that  sharp  dissociation  at  this  particular  point 
which  is  necessary  for  the  production  of  the  great  excess  of  free 
carbon  needed  to  furnish  light. 

Before  these  investigations  were  made,  a  different  answer  was 
given  to  the  question  why  the  flame  of  pure  illuminating  gas  con- 
tains free  carbon  and  is  luminous.  It  was  said  that  hydrogen  was 
more  easily  burned  than  carbon,  and  therefore  the  latter  was  left 
free,  to  be  burned  later.  It  is  true  that  gaseous 
hydrogen  burns  more  easily  than  solid  carbon  e.g., 
charcoal.  But  in  ethylene,  both  elements  are  equally 
gaseous  and  the  explanation  is  faulty.  Smithells 
(1892)  demonstrated  the  falsity  of  this  explanation 
by  devising  a  cone-separator  (Fig.  135).  The  air 
and  ethylene  or  other  gas  are  admitted  in  propor- 
tions which  can  be  varied,  and  the  mixture  burns  at 
the  top  of  the  wider  tube.  As  the  quantity  of  air  is 
increased,  however,  the  speed  with  which  an  explosion- 
flame  would  pass  through  it  becomes  greater,  and 
finally  the  inner  cone  passes  down  and  rests  upon  the 
mouth  of  the  narrow  tube  through  which  the  mixture 
of  gases  is  issuing  more  rapidly.  A  preliminary  com- 
bustion takes  place  in  the  blue  cone,  while  the  final  con-  FlG-  135-  ' 
version  of  the  whole  material  into  carbon  dioxide  and  water  is 
completed  in  the  outer  mantle.  This  remains  at  the  top  of  the 
wider  tube,  where  alone  the  necessary  air  can  be  obtained.  By 
means  of  a  side  tube  (not  shown)  he  withdrew  the  inter-conal  gas 
and  found  that,  while  all  of  the  carbon  was  burned  by  the  inner 
cone  as  far  as  carbon  monoxide  CO,  most  of  the  hydrogen  was  still 
entirely  uncombined.  This  was  the  case,  not  only  with  illuminating- 
gas,  which  initially  contains  much  free  hydrogen,  but  also  when  the 
flame  was  fed  with  pure  methane.  The  change  in  the  inner  cone  of 
the  Bunsen  flame  consists,  therefore,  mainly  in  the  burning  of  all 
the  hydrocarbons  to  carbon  monoxide,  with  liberation  of  the  hydro- 
gen. In  the  outer  cone,  it  is  practically  a  burning  of  water  gas 
that  is  taking  place. 


602  INORGANIC  CHEMISTRY 

Exercises.  —  1.   Write  a  graphic  formula  for  hexane. 

2.  Write  an  equation  for  the  formation  of  aluminium  carbide 
(p.  589). 

3.  Make  a  section  showing  the  shape  of  the  flame  produced  by 
burning  hydrogen  gas  when  the  latter  issues  from  a  circular  opening. 

4.  Name  the  radicals  C5Hn,  C5Hi0,  CsHg,  CieHsi,  and  the  sub- 
stances CsHiiCl,  C5HiiHSO4. 


CHAPTER  XXX 

THE   CARBOHYDRATES,    ORGANIC  ACIDS,   ALCOHOLS,    SOAP, 
COLLOIDS,   FOODS 

A  PLANT  takes  carbon  dioxide  from  the  air  and  water  from  the 
ground  and,  using  the  energy  of  sunlight,  converts  them  into  a 
growing  framework  of  cellulose  (CoKioO^x  and,  as  we  have  seen 
(p.  580),  into  starch  (CeHioOs)^  which  it  stores  in  the  cells.  The 
cellulose  of  certain  plants  furnishes  us  with  cotton,  linen,  jute, 
and  paper.  The  starch  of  wheat,  oats,  maize  (corn),  and  potatoes 
is  one  of  the  chief  foodstuffs  they  contain.  The  plant,  when  dead 
and  buried,  changes  into  coal.  The  fresh  wood,  when  distilled, 
supplies  wood  spirit  (methyl  alcohol)  and  other  useful  substances, 
and  the  residue  is  the  valuable  charcoal.  Furthermore,  from  starch 
we  can  readily  make  sugar,  alcohol,  and  other  familiar  materials. 
Cellulose,  starch,  and  the  sugars  (e.g.,  cane-sugar  C^H^On)  contain 
oxygen  and  hydrogen  in  the  same  proportions  in  which  they  are 
present  in  water,  2H  :  1O.  They  might  be  considered  hydrates  of 
carbon,  and  so  they  are  called  the  carbohydrates.  The  foregoing 
brief  summary  shows  that  the  carbohydrates  introduce  us  to  a  much 
greater  variety  of  interesting  organic  compounds  than  does  petroleum. 

THE  CARBOHYDRATES 

Cellulose  (C&Hi0O5)x  and  Paper.  —  The  wall  of  each  cell,  and 
therefore  the  whole  framework  of  a  plant,  is  made  of  cellulose. 
Linen  and  cotton  are  pure  cellulose.  The  walls  of  the  cells  are 
usually  more  or  less  thickened  by  a  substance  called  lignin,  which 
has  much  the  same  composition,  but  different  chemical  behavior. 
The  best  paper  is  made  of  cotton  or  linen  (rag-paper).  Cheaper 
kinds  are  prepared  by  cutting  wood,  such  as  spruce  or  pine,  into 
chips  and  treating  ("cooking")  them  with  a  solution  of  calcium 
bisulphite  Ca(HSOs)2.  This  process  decomposes  the  lignin,  and 
converts  it  into  soluble  materials.  The  sulphite  liquor  is  then 
run  off,  and  the  pulpy  material  is  washed,  beaten  with  water  to 
reduce  it  to  minute  shreds,  and  bleached  with  very  dilute  chlorine- 
water.  The  pure  cellulose,  now  paper-pulp,  is  suspended  in 

603 


604  INORGANIC  CHEMISTRY 

water,  spread  on  screens,  pressed,  and  dried.  During  the  process, 
other  substances  are  added.  Thus,  size  (gelatine  or  rosin  and 
alum,  see  Sizing)  prevents  the  ink  from  running;  pulverized  cal- 
cium sulphate  (gypsum),  clay,  and  other  white  solids  ("  loading") 
give  body  to  the  paper  and  permit  the  production  of  a  smooth  sur- 
face by  rolling  ("calendering").  Dyestuffs  can  be  added  to  give 
special  tints.  Filter  paper  is  pure  cellulose. 

The  sulphite  lyes,  after  neutralization  of  the  acids,  and  other 
treatment,  can  be  fermented  with  yeast  to  give  alcohol  (see  p.  607). 
Pure  cellulose  gives  no  sugar,  so  that  the  alcohol  comes  from  decom- 
position products  of  the  lignin. 


Starch  (C^H^O^y.  —  Starch  consists  of  little  colorless  granules 
of  various  rounded  shapes  (Fig.  136),  which  are  easily  seen  under 
the  microscope.  These  granules  are  massed  in  large  quantities 

in  the  ears  of  wheat  and  oats,  in  the  tubers 
of  potatoes,  in  the  grains  of  maize  (corn), 
and  in  peas  and  beans.  Even  in  the  leaves 
they  can  be  seen.  Starch  is  recognized  by 
the  iodine  test  (p.  276),  turning  deep  blue 
with  a  trace  of  free  iodine. 
**y  The  treatment  of  wheat  flour,  which 

is  three-fourths  starch,  by  washing  out 

the  starch  through  porous  cloth  with  water,  has  already  been  de- 
scribed (p.  5).  Starch  is  made  from  maize  in  America  and  from 
potatoes  in  Europe,  by  washing  the  flour  on  sieves. 

Starch  is  not  soluble  in  water.  If  boiled  with  water,  however, 
the  granules  swell  and  break  and  the  starch  is  diffused  through 
the  water,  giving  a  clear  liquid.  If  too  much  water  is  not  used,  the 
liquid  when  cold  sets  as  a  jelly.  While  the  liquid  is  hot,  much  of 
the  starch  will  pass  through  a  filter  along  with  the  water.  Such 
a  liquid  is  called  a  colloidal  suspension.  Suspensions  like  this  are 
constantly  met  with  in  using  complex  organic  compounds  like  jellies, 
glues,  soaps,  and  dyes.  Even  insoluble  inorganic  materials,  like 
gold,  give  such  suspensions  (see  p.  621). 

The  colloidal  suspension  of  starch  turns  blue  when  a  solution 
containing  free  iodine  is  added  to  it.  It  is  used  in  the  laundry  for 
stiffening  white  goods.  Glucose  is  manufactured  from  it. 

Glucose  C6Hi2O6,  a  Sugar,  from  Starch.  —  When  starch  is 
boiled  with  water,  to  which  a  few  drops  of  an  acid  (contact  agent), 


CARBOHYDRATES  605 

such  as  hydrochloric  acid,  have  been  added,  the  liquid,  after  neutrali- 
zation of  the  acid,  is  found  to  be  sweet.  One  of  the  sugars,  glucose 
CeH^Oe,  can  be  obtained  in  crystals  after  evaporation.  The  crystals 
form  "brewers'  glucose"  and  the  syrup  produced  by  concentration 
is  corn  syrup  (if  maize  is  the  source  of  the  starch).  The  latter, 
although  less  sweet  than  ordinary  sugar,  is  much  less  expensive 
and  is  used  in  making  preserves  and  cheap  candy. 

The  molecular  weight  of  starch  is  at  least  as  large  as  (C6Hi005)2oo. 
The  formula  (CeHioOs)^  shows  the  composition.  The  water,  in 
presence  of  a  little  acid,  decomposes  the  molecules  and  combines 
with  the  material.  First,  dextrin  (used  as  paste  or  mucilage)  is 
formed  and  this  breaks  up  into  glucose.  The  action  is  an  hydrolysis  . 

(C6Hio05)tf  +  2/H20  ->  2/C6H]206. 

Glucose  is  known  also  as  dextrose  and  as  grape  sugar.  The 
crystalline  granules  in  raisins  (dried  grapes)  are  composed  mainly 
of  it.  When  pure,  it  is  almost  colorless.  It  reduces  cupric  hydrox- 
ide, in  Fehling's  solution  (q.v.),  to  cuprous  oxide. 

Corn  syrup  contains  30-40  per  cent  of  unchanged  dextrin,  40-50 
per  cent  of  glucose,  and  the  rest  water. 

The  Sugars.  —  The  common  sugars  may  be  divided  into  the 
monosaccharides,  usually  with  the  formula  CeH^Oe,  and  the  disaccha- 
rides,  usually  C^H^On.  Of  these,  the  following  will  be  referred  to 
in  what  follows: 


Monosaccharides:  Glucose    (grape  sugar  or  dextrose) 
Fructose  (fruit  sugar  or  levulose) 

Disaccharides  :        Sucrose    (cane  sugar,  beet  sugar) 
Maltose    (action  of  malt  on  starch) 
Lactose    (milk  sugar,  in  animals  only)  C^H^Ou. 

Sucrose,  or  Cane  Sugar  CwHmOu..  —  Plants,  such  as  the  sugar- 
cane and  beet,  besides  forming  cellulose  and  starch,  produce  excep- 
tional amounts  of  sucrose,  or  table  sugar.  The  sap  of  the  sugar 
maple  contains  much  of  it. 

Cane  sugar  is  extracted  by  crushing  the  stalks  between  rollers, 
and  evaporating  the  expressed  liquid  (18  per  cent  sugar)  in  closed 
pans.  A  partial  vacuum  is  maintained  so  that  the  solution  may 
boil  at  a  low  temperature  (65°  to  begin  with)  and  none  of  the  sugar 
be  decomposed.  When  the  syrup  cools,  the  sucrose  appears  in 


606  INORGANIC  CHEMISTRY 

brown-colored  crystals.  The  mother-liquor  is  called  molasses. 
In  the  refinery,  the  sugar  is  redissolved,  the  solution  is  poured  through 
a  column  of  charcoal,  which  takes  out  the  coloring  matter,  and  the 
liquid  is  once  more  allowed  to  deposit  crystals.  Pure  cane-sugar 
has  a  faint  yellow  lint,  and  a  small  amount  of  ultramarine  (q.v.)  is 
added  to  give  that  whiteness  which  is  popularly  connected  with 
purity  in  sugar. 

Sugar  beets  (16  per  cent  or  more  sugar)  are  sliced  and  steeped 
in  water.  The  extract  contains  a  gummy  material  in  colloidal 
suspension.  This  is  coagulated  and  precipitated  by  adding  slaked 
lime  (calcium  hydroxide)  Ca(OH)2  suspended  in  water,  and  boil- 
ing. After  separation  of  the  clear  liquid,  carbon  dioxide  is  passed 
through  it  to  precipitate  the  excess  of  lime  as  carbonate  CaC03. 
The  solution  is  then  decolorized  with  charcoal  and  evaporated  to 
crystallization. 

As  regards  properties,  sucrose  crystallizes  in  four-sided  prisms 
(rock-candy,  Fig.  73,  p.  173)  and  melts  at  160°.  When  heated  to 
200-210°  it  partially  decomposes,  leaving  a  soluble,  brown,  mixed 
material,  caramel,  used  in  coloring  whisky  and  soups.  Sucrose  does 
not  reduce  Fehling's  solution. 

When  boiled  with  water  containing  a  trace  of  acid  (contact 
agent),  sucrose  is  hydrolyzed,  giving  a  mixture  of  the  two  mono- 
saccharides,  glucose  and  fructose: 

Ci2H22On  +  H20  ->  C6H12O6  +  C6Hi2O6. 

The  process  is  called  inversion,  and  the  mixture  invert-sugar.  It  is 
found  in  many  sweet  fruits  and  in  honey.  Each  sugar  interferes 
with  the  crystallization  of  the  other,  by  lowering  the  freezing-point 
(p.  335),  and  so  invert-sugar  is  added  in  making  fondant  candy  and 
candy  that  is  to  be  pulled,  both  of  which  must  remain  soft  for  some 
time.  '  Icing  for  cakes  has  to  some  extent  this  property,  and  is  made 
by  adding  acid  substances  to  sugar,  such  as  vinegar,  lemon  juice,  or 
cream  of  tartar. 

The  action  of  the  acid  in  inversion  is  catalytic,  and  the  rate  of  this 
more  or  less  leisurely  chemical  change  depends  upon  the  concentra- 
tion of  the  hydrogen  ions.  It  therefore  furnishes  one  means  of 
comparing  acids  as  regards  their  chemical  activity,  and  has  the 
special  advantage  that  the  acid  is  not  consumed  during  the  process 
(cf.  p.  295),  but  remains  of  constant  concentration  throughout  the 
whole  time. 


FERMENTATION  607 

Enzymes.  —  These  are  active  organic  compounds  contained  in 
certain  vegetable  organisms.  The  organisms  may  be  divided  into 
three  classes,  each  secreting  different  enzymes,  which  confine  them- 
selves for  the  most  part  to  special  kinds  of  chemical  change.  (1) 
The  molds,  when  grown  in  sugar  solution  or  beef  extract,  or  other 
nutritive  solutions,  produce  decompositions  known  collectively  as 
putrefaction.  (2)  Certain  bacteria  promote  the  oxidation  of  alcohol 
to  acetic  acid  (see  p.  609).  Some  also  decompose  sugar,  furnishing 
butyric  or  lactic  acid  as  one  of  the  products.  (3)  The  yeasts  (sac- 
char  omycetes),  decompose  sugars  into  alcohol  and  carbon  dioxide. 
They  consist  of  microscopic  cells,  which,  while  multiplying,  secrete 
within  each  cell  two  very  active,  soluble  substances.  These  are 
zymase  and  sucrase  (invertase),  which  belong  to  a  class  of  organic 
substances  called  enzymes.  Sucrase  means  an  enzyme  that  splits 
sugar.  Enzymes  produce  remarkable  chemical  changes  by  their 
mere  presence  (contact  actions). 

Enzymes  which  produce  still  other  kinds  of  chemical  change  are 
found  in  the  body,  such  as  pepsin  which  hydrolyzes  proteins  in  the 
stomach. 

Alcoholic  Fermentation.  —  When  some  yeast  (S.    cerevisice), 
which  is  a  mass  of  the  living  plants,  is  added  to  a  solution  of  glucose 
at  about  30°,  the  small  amount  of  zymase  gradually  decomposes 
the  sugar.     Bubbles  of  carbon  dioxide  soon  begin  to 
rise,   and  may  be    tested   (p.   501)   with   limewater 
(Fig.  137).     At  the  same  time,  alcohol  (ethyl  alcohol 
C2H5OH)  accumulates  in  the  liquid: 

C6Hi206  ->  2C02  T  +  2C2H5OH. 

Yeast  will  ferment  fructose  CeH^Oe  with  the  same    Z_ _^ 

result,  but  more  slowly,  so  that,  when  placed  in  in- 
vert sugar,  it  decomposes  the  glucose  first  and  the 
fructose  afterwards.  FIQ  13? 

Zymase  does  not  act  upon  sucrose  (table  sugar), 
but  sucrase  hydrolyzes  the  sucrose  in  the  same  way  as  does  a  dilute 
acid,  giving  invert-sugar.     The  latter  is  then  decomposed  by  the 
zymase,  and  so  cane-sugar  in  solution  is  fermented  by  yeast  into 
alcohol  and  carbon  dioxide,  just  as  is  glucose,  only  more  slowly. 

In  making  wine,  the  glucose  contained  in  the  grape  juice  is  fer- 
mented by  a  species  of  yeast  found  on  the  skins  of  the  grapes  (S. 
ellipsoideus).  The  wine  is  allowed  to  stand,  after  fermentation,  m> 


608  INORGANIC  CHEMISTRY 

til  it  has  deposited  a  considerable  crust  of  material  known  as  argol, 
which  consists  mainly  of  potassium-hydrogen  tartrate  KHC4H4O6, 
cream  of  tartar.  The  concentration  of  the  sugar  in  the  grape- 
juice  being  small,  the  quantity  of  alcohol  contained  in  the  product  is 
not  very  great.  By  distillation  of  wine,  brandy,  containing  a  much 
larger  proportion  of  alcohol,  is  made.  The  special  flavors  of  wines 
and  brandies  depend  upon  materials,  other  than  sugar,  originally 
contained  in  the  fermented  liquid,  upon  by-products  of  the  fermenta- 
tion, and  upon  materials  which  arise  by  slow  chemical  changes 
while  the  liquor  is  stored. 

Commercial  alcohol  is  not  made  from  sugar,  but  from  the  starch 
of  potatoes  or  maize.  When  barley  is  allowed  to  sprout,  an  enzyme, 
amylase  (meaning  starch-splitting  enzyme)  or  diastase,  is  formed  in 
the  ears.  The  whole  material  is  dried,  and  is  then  called  malt. 
When  this  is  mixed  with  starch  and  water,  the  amylase  hydrolyzes 
the  starch  to  maltose  C^H^Ou  (p.  605).  The  latter  is  then  further 
hydrolyzed  by  yeast  to  form  glucose  CeH^Oe,  and  this  is  decomposed 
by  the  zymase  into  alcohol  and  carbon  dioxide. 

Whisky  (about  50  per  cent  alcohol)  is  made  by  treating  the 
starch  of  rye,  maize,  or  barley  in  the  same  way,  with  subsequent 
distillation  (see  below)  to  separate  the  alcohol  (whisky).  Beer  is 
made  similarly  from  various  kinds  of  grain,  especially  barley,  but 
the  fermented  liquor  is  not  distilled.  Aside  from  the  alcohol  and 
carbon  dioxide,  considerable  quantities  of  other  substances  extracted 
from  the  grain  remain  in  the  solution  and  form  the  so-called  "ex- 
tract," which  varies  in  kind  and  quantity  in  different  varieties  of 
beer. 

When  used  in  bread-making,  the  yeast  acts  upon  a  trace  of  sugar 
in  the  flour,  and  the  carbon  dioxide  gas  evolved  raises  the  bread. 

Ethyl  Alcohol  C^H^OH.  —  Common  alcohol  is  related  to 
ethane  CeH6,  having  an  hydroxyl  group  in  place  of  one  unit  of 
hydrogen.  Hence  its  name. 

Ethyl  alcohol  boils  at  78.3°  and  so,  when  the  fermented  liquor 
is  distilled  (rectified)  it  is  almost  pure  alcohol  that  comes  off.  Com- 
mercial alcohol  contains  95  per  cent  by  volume  (in  Great  Britain, 
90  per  cent).  Absolute  alcohol  cannot  be  made  by  distillation 
alone  (see  below).  It  is  prepared  by  adding  quicklime,  which 
combines  with  the  water,  and  redistilling  the  liquid. 

Alcohol  mixes  with  water  in  all  proportions.  In  dilute,  aqueous 
solution  it  is  not  ionized,  and  does  not  interact  with  acids,  bases, 


FERMENTATION  609 

or  salts.  It  is,  however,  easily  oxidized  to  acetic  acid.  When 
water  is  absent,  it  interacts  with  acids  slowly  (see  p.  617). 

Alcohol  is  used  as  a  solvent  for  the  resins  employed  in  making 
varnishes  for  wood  and  lacquers  for  metal. 

On  account  of  the  high  duty  on  95  per  cent  alcohol  ($2.11  per 
gallon  in  the  U.  S.  and  24/6  in  Gt.  Britain),  denatured  alcohol 
(methylated  spirit),  which  is  free  of  duty,  is  employed  for  indus- 
trial purposes.  The  alcohol  (cost  about  22  cents  per  gal.  hi  the 
U.  S.  and  1/6  in  Britain)  is  mixed  with  offensive  or  poisonous  ma- 
terials, which  prevent  its  consumption  as  a  beverage,  without  inter- 
fering with  other  uses.  Wood  spirit  and  gasoline  are  often  employed. 

Distillation  of  Ethyl  Alcohol.  —  Mixtures  of  two  liquids, 
when  distilled,  behave  in  one  of  three  ways.  Two  of  these  have 
been  described  already  (p.  212),  and  alcohol  (b.-p.  78.30°)  and 
water  (b.-p.  100°)  illustrate  the  third.  In  this  case  the  vapor  tension 
of  a  certain  mixture  is  higher  than  that  of  any  other  mixture  and 
higher  than  that  of  either  component  separately.  This  special 
mixture  has,  therefore,  a  lower  boiling-point  than  any  other.  In 
the  present  instance  this  mixture  contains  95.57  per  cent  of  alcohol 
and  4.43  per  cent  of  water  and  boils  at  78.15°.  When  the  fermented 
liquid,  with  its  large  percentage  of  water,  is  distilled,  the  alcohol  all 
tends  to  pass  off  first,  in  association  with  that  part  of  the  water 
required  to  constitute  the  mixture  of  minimum  boiling-point.  Re- 
peated distillation  simply  eliminates  more  completely  the  excess 
of  water  beyond  this  amount  (viz.,  4.43  per  cent),  by  leaving  it  in  the 
residues. 

Acetic  Acid  HCO2CHS.  —  This  is  the  sour  substance  in  vinegar, 
and  has  many  industrial  applications.  Vinegar  is  made  by  oxi- 
dizing alcohol  with  atmospheric  oxygen,  using  an  enzyme  secreted 
by  bacterium  aceti  (mother  of  vinegar)  as  contact  agent.  Oxygen 
alone  does  not  affect  alcohol  in  the  cold.  Dilute  alcohol  from  any 
source,  such  as  fermented  apple  juice  (hard  cider),  is  allowed  to 
trickle  over  shavings  in  a  barrel.  Holes  admit  air,  and  the  shavings 
are  inoculated  in  advance  by  wetting  with  vinegar: 

HOCH2CH3  +  O2->HC02CH3  +  H2O. 

The  issuing  liquid  contains  5-15  per  cent  of  acetic  acid,  which  can 
be  purified  by  fractional  distillation  to  separate  the  water.  It 
boils  at  118°  and  freezes  at  16.7°.  Although  four  atoms  of  hydro- 


610  INORGANIC  CHEMISTRY 

gen  are  contained  in  its  molecule,  but  one  of  these  is  replaceable 
by  metals.  This  fact  is  recognized  in  the  reaction  formula  of 
the  acid,  HC2H302,  or  HC02CH3.  It  is  a  weak,  monobasic  acid: 
HC02CH3  +±  H+  +  CO2CHr. 

WOOD,  CHARCOAL,  COAL,  COKE 

Destructive  Distillation  of  Wood.  Charcoal.  — *  Dry  wood 
is  distilled  in  iron  retorts,  and  the  vapors  coming  off  are  led  through 
a  condenser  to  separate  the  liquids  from  the  gases.  The  cellulose, 
lignin,  moisture,  and  resinous  material  are  decomposed  or  volatilized, 
and  only  charcoal  remains.  The  gases,  consisting  mainly  of  hydrogen, 
methane  CH4,  ethane  C2He,  ethylene  C2H4,  and  carbon  monoxide 
CO,  are  employed,  on  account  of  their  combustibility,  as  fuel  in  the 
distillation  itself.  The  fluids  form  a  complex  mixture  containing 
large  quantities  of  water,  methyl  alcohol  CH3OH  (wood  spirit), 
acetic  acid,*  acetone  (CH3)2CO,  and  tar.  The  liquids  can  be  sepa- 
rated. The  methyl  alcohol  (wood  spirit)  is  used  in  varnish  making. 
The  acetone  has  several  uses  (e.g.,  p.  593). 

Wood  charcoal  exhibits  the  cellular  structure  of  the  material 
from  which  it  was  made,  and  is  therefore  highly  porous  and  has 
an  enormous  internal  surface.  When  the  charcoal  is  burned,  the 
mineral  constituents  of  the  wood  appear  in  the  ash.  This  is  com- 
posed of  the  carbonates  of  the  metallic  elements  present.  For 
certain  purposes,  charcoals,  made  in  the  same  fashion  as  the  above 
from  bones  and  from  blood,  find  wide  application.  The  former, 
called  bone  black,  contains  much  calcium  phosphate  (p.  548).  In 
the  old  method  of  making  charcoal,  which  is  still  practiced,  the 
wood  was  piled  up,  covered  with  turf,  and  set  on  fire.  All  the  valua- 
ble volatile  products  were  lost,  as  well  as  part  of  the  charcoal  itself. 

In  the  laboratory,  pure  carbon  is  made  by  dissolving  sugar  in 
little  water  and  adding  concentrated  sulphuric  acid  as  a  dehydrating 
agent:  Ci2H22On  — >  12C  +  11H2O.  The  black  mass  is  washed 
with  water  until  free  from  the  acid.  The  sugar  is  purified  from 
mineral  matter,  before  use,  by  crystallization  from  water. 

Properties  of  Charcoal.  —  Charcoal  is  amorphous  carbon  (den- 
sity variable,  up  to  1.8).  It  exhibits  certain  properties  which  are 

*  The  dry  distillation  of  bones  (see  below),  on  the  other  hand,  and  of  animal 
matter  (p.  520)  in  general,  gives  alkaline  liquids,  because  of  the  ammonia  that  is 
formed. 


WOOD,   CHARCOAL,   COAL,  COKE  611. 

not  shared  by  other  forms  of  carbon.  For  example,  it  can  take 
up  large  quantities  of  many  gases.  Boxwood  charcoal  will  in  this 
way  absorb  ninety  times  its  own  volume  of  ammonia,  fifty-five  vol- 
umes of  hydrogen  sulphide  or  nine  volumes  of  oxygen.  Freshly 
made  dogwood  charcoal  (used  in  making  the  best  gunpowder), 
when  pulverized  immediately  after  its  preparation,  often  catches  fire 
spontaneously  on  account  of  the  heat  liberated  by  the  condensation 
of  oxygen.  It  is  therefore  set  aside  for  two  weeks,  to  permit  the  slow 
absorption  of  moisture  and  air.  The  absorbed  gases  may  be  removed 
unchanged  by  heating  the  charcoal  in  a  vacuum.  The  phenomenon, 
described  as  adsorption,  is  caused  by  the  adhesion  of  the  gases  to  the 
very  extensive  surface  (due  to  porosity)  which  the  charcoal  possesses. 
Glass  and  all  other  solids  show  the  same  property,  though  in  a  smaller 
degree  (p.  147).  Solid  and  liquid  bodies  are  also  in  many  cases  taken 
up  by  charcoal  in  a  similar  fashion.  Thus,  organic  dyes,  such  as 
indigo,  litmus,  and  cochineal,  and  natural  coloring  matters  (see 
sugar  refining,  p.  606),  which  are  all  more  or  less  colloidal  in  nature, 
are  removed  when  the  liquid  is  shaken  with,  or  poured  through 
pulverized  charcoal.  The  organic  materials  dissolved  in  the  drinking 
water  also  undergo  adsorption  in  charcoal,  but  the  charcoal  soon  be- 
comes inactive.  Charcoal  is  likewise  used  in  reducing  ores,  and 
as  a  smokeless  fuel. 

Coal.  —  When  vegetable  matter  decomposes,  without  heating, 
and  while  covered  with  sand  or  clay  so  that  air  is  excluded;  water 
and  hydrocarbons  are  liberated,  and  the  products  are  peat,  or  bitumi- 
nous coal,  or  anthracite. 

We  are  concerned  mainly  with  the  products  obtained  by  dis- 
tilling coal,  to  get  coal  gas  and  coke,  and  with  its  use  as  fuel.  To 
determine  its  suitability  for  various  purposes,  the  coal  is  analyzed, 
and  its  heating  power  is  measured. 

In  coal  analysis,  the  air-dried  material  is  used.  The  water  is 
determined  by  heating  1  g.  at  105°  for  1  hour.  Much  water  lowers 
the  fuel  value,  because  heat  is  wasted  in  vaporizing  it,  and  in  de- 
composing it  (cf.  p.  577).  After  reweighing  the  sample,  the  coal  is 
heated  with  the  Bunsen  flame  in  a  covered  crucible  to  drive  off 
the  volatile  matter.  After  weighing  again,  air  is  admitted,  and 
strong  heat  is  applied  to  burn  up  the  fixed  carbon  (coke).  The 
residue  is  the  ash.  In  the  following  table  the  proportions  are  com- 
pared with  seasoned  wood  on  the  one  hand,  and  with  charcoal  and 
coke  on  the  other. 


612 


INORGANIC  CHEMISTRY 


The  calorific  power  of  a  coal  determines  largely  its  value  for 
heating.  A  sample  (about  1  g.)  is  burned  in  a  bomb  calorimeter 
(p.  98).  The  rise  in  temperature  of  a  known  weight  of  surround- 
ing water  gives  the  number  of  calories.  The  coal  is  set  on  fire  by 
a  wire  heated  electrically.  Engineers  use  the  number  of  British 
Thermal  Units  (1  B.T.U.  =  heat  required  to  raise  1  Ib.  of  water  1°  F.) 
developed  by  1  pound  of  coal.  The  number  of  B.T.U.  =  1.8  X 
number  of  calories  per  grain  of  coal. 

Bituminous  coals  give  much,  and  widely  varying  amounts  of 
volatile  matter,  while  anthracite  gives  very  little.  The  ash  is  the 
mineral  matter  of  the  original  plants,  with  rock  material  in  many 
specimens.  For  coal  gas,  and  even  for  coke,  a  coal  high  in  volatile 
matter  is  chosen.  For  water  gas  (p.  577)  anthracite  or  coke  is 
employed. 


Water. 

Vola- 
tile 
matter. 

Fixed 
car- 
bon. 

Ash. 

Sul- 
phur. 

Gal.  per 
lg. 

Wood      

20.0 

49.0 

30.0 

1.0 

3,100 

Peat    

20.0 

51.6 

25.0 

3.2 

0  2 

4,270 

Bituminous  

1.3 

36.7 

53.5 

8.5 

1  7 

7,800 

Semi-bit  um  i  no  us 

4  0 

16  0 

68  5 

11 

0  5 

7  510 

Anthracite 

3  0 

5  6 

80  5 

10  9 

0  8 

8000 

Charcoal 

3  2 

4  2 

90  7 

1  7 

7  580 

Coke 

2  5 

1  3 

86  3 

12  4 

1  3 

7  770 

Petroleum     . 

11,000 

i  '  If  the  heat  of  combustion  of  a  coal  is  known,  the  amount  of 
steam  it  should  furnish  can  be  calculated.  It  takes  100  cal.  to 
raise  1  g.  of  water  from  0°  C.  to  100°  C.  and  540  cal.  more  to  con- 
vert it  into  steam.  If  the  quantity  of  steam  is  too  small,  the  furnace, 
draft,  or  firing  is  defective.  Too  much  draft,  for  example,  merely 
adds  additional,  useless  air  to  be  heated.  If  the  flue  gas,  when 
analyzed,  contains  only  3  per  cent  of  carbon  dioxide,  instead  of  the 
normal  12  per  cent,  then  for  every  ton  of  coal  burned,  52  tons  of 
unnecessary  air  were  raised  to  the  temperature  of  the  furnace.  Tests 
of  this  kind  can  control  the  efficiency  of  every  device  in  a  modern 
factory,  and  they  ought  to  be  in  universal  use. 

Coal  Gas.  — The  gas  plant  (Fig.  138)  includes:  (1)  The  fire- 
brick retorts  in  which  the  coal  is  heated  to  1300°,  (2)  the  hydraulic 
main,  a  wide  iron  pipe  above  them  in  which  the  tar  collects,  (3) 
the  condenser  and  wash  box  for  cooling  and  condensing  oils,  (4) 


WOOD,   CHARCOAL,   COAL,   COKE 


613 


the  scrubbers  where  the  ammonia  is  taken  out  by  water,  (5)  the 
purifier  where  hydrogen  sulphide  is  absorbed  by  hydrated  ferric 
oxide  and  (6)  the  holder  where  the  gas  collects. 


FIG.  138. 

One  short  ton  (2000  Ibs.)  of  the  bituminous  coal  in  the  above 
table  gave:  Gas  10,500  cu.  ft.  with  13  candle  power,  coke  1325  Ibs., 
ammonia  5  Ibs.  (=  20  Ibs.  (NH4)2SO4  worth  $60  per  ton),  and  tar 
12  gallons.  The  components  of  the  gas  were:  Illuminants  3.8, 
heating  gases  90.2,  impurities  6.0.  Calorific  power  of  gas  610  B.T.U. 
per  cu.  ft.;  sp.  gr.  (air  =  1)  0.43. 

The  tar  is  frequently  distilled  fractionally  and  yields:  Benzene 
CeHe,  from  which  aniline  and  many  dyes  and  drugs  are  prepared; 
naphthalene  CioHs,  sold  as  moth-balls,  and  the  starting  point  for 
synthetic  indigo;  anthracene  CnHio,  from  which  valuable  dyes 
such  as  alizarin  and  indanthrene  are  made;  phenol  or  carbolic 
acid  (p.  527),  a  disinfectant,  and  other  useful  substances.  A  rougher 
separation  yields  tar  and  pitch,  for  road-making,  preserving  timber, 
and  waterproofing  roofs. 

Coke.  —  The  beehive  coke  oven  is  a  brick  structure  shaped  like 
a  beehive,  with  an  additional  opening  at  the  top.  The  coal  which 
fills  it  burns  with  a  limited  supply  of  air.  All  the  vapors  and  gas 
burn  at  the  upper  opening,  and  the  ammonia  and  tar  and  com- 
bustible gas  are  therefore  wasted  (cf.  p.  516). 

The  by-product  coke  oven  is  a  good  deal  like  a  gas  plant.  The 
chief  difference  is  that  the  heating  is  arranged  so  as  to  decompose 
as  much  of  the  volatile  matter  as  possible,  and  cause  it  to  leave  its 
carbon  in  the  retort.  The  gas  is  therefore  poor  in  illuminants,  but 
excellent  as  fuel.  The  ammonia  and  tar  are  diminished  in  amount, 
but  still  valuable  products.  The  yield  of  coke  is  about  73  per  cent 
of  the  original  coal,  against  66  per  cent  from  the  beehive  oven. 

Burning  coke  gives  a  higher  temperature  than  does  coal,  be- 


614  INORGANIC  CHEMISTRY 

cause  no  heat  is  used  in  vaporizing  moisture  and  volatile  matter. 
For  the  same  reason,  it  burns  without  flame.  Because  of  these 
and  other  properties,  it  is  employed  in  immense  quantities  in  re- 
ducing ores  of  iron  in  the  blast  furnace,  as  well  as  for  many  other 
purposes. 

ORGANIC  ACIDS  AND  SALTS 

Thus  far,  one  acid,  acetic  acid,  and  two  alcohols,  methyl  and 
ethyl  alcohol,  have  been  mentioned.  But  there  are  whole  series 
of  organic  acids  and  alcohols,  corresponding  to  the  series  of  hydro- 
carbons. 

Organic  Acids  and  Their  Salts.  —  The  general  formula  of  the 
saturated  series  of  monobasic  acids  is  H(CO2CnH2n+i).  Thus: 

Formic  acid  (n  =  0),  H(CO2H).  Palmitic  acid  (n  =  15),  H(CO2Ci5H3i). 
Acetic  acid  (n  =  1),  H(CO2CH3).  Stearic  acid  (n  =  17),  HCCOaCnHj. 
Butyric  acid  (n  =  3),  H(CO2C3H7). 

Formic  (Lat.  formica,  an  ant)  acid  is  secreted  by  red  ants.  Formic 
(b.-p.  100.1°),  acetic,  and  butyric  acids  are  liquids.  Palmitic  and 
stearic  acids  are  solids,  and  are  mixed  with  paraffin  in  making  candles. 

Acids  containing  less  hydrogen  are  unsaturated.  Thus,  oleic 
acid  (n  =  17)  is  H(CO2Ci7H33). 

The  acids  with  large  molecular  weight  are  insoluble  in  water. 
All  the  acids,  however,  react  with  sodium  hydroxide  solution,  giving 
soluble  salts.  Thus,  palmitic  acid  gives  sodium  palmitate: 

NaOH  +  H(CO2Ci5H3i)  <=*  H2O  +  Na(CO2Ci5H3i). 

Other  salts  are  sodium  formate  (p.  577)  Na(C02H),  sodium  acetate 
Na(CO2CH3),  sodium  stearate  Na(CO2Ci7H35),  sodium  oleate  Na 
(CO^ffHss).  Common  soap  is  a  mixture  of  the  last  two  salts  with 
sodium  palmitate. 

Later,  in  discussing  fats  and  soap,  it  will  be  convenient  to  abbre- 
viate the  formulae.  A  monobasic  acid  will  be  indicated  by  the 
formula  HCO2R  and  a  salt  by  NaCO2R  or  Ca(CO2R)2,  where  R 
stands  for  a  hydrocarbon  radical  or  group  of  atoms,  such  as  CnH2n+i 
In  organic  chemistry  a  radical  is  not  always  able  to  form  an  ion. 
Here  the  ion  is  CO2R~. 

Formic  Acid  HCO^H.  —  The  removal  of  water  from  formic 
acid  produces  carbon  monoxide  (p.  577).  Although  we  cannot 


ORGANIC  ACIDS  AND  SALTS  615 

reverse  the  process  and  cause  carbon  monoxide  to  combine  with 
water,  yet  by  passing  carbon  monoxide  over  hot  sodium  hydroxide, 
we  obtain  sodium  formate,  from  which  formic  acid  may  be  liber- 
ated by  double  decomposition  with  another  acid:  CO  +  NaOH  —  > 
NaC02H. 

This  acid  is  secreted  by  red  ants,  and  is  found  in  stinging  nettles. 
It  is  a  liquid  boiling  at  100.1°  and  freezing  at  8.6°.  Although  one  of 
the  weaker  acids,  it  is  much  more  active  than  acetic  acid.  The 
molecule  contains  two  atoms  of  hydrogen,  but  the  acid  is,  in  fact, 
monobasic.  The  structural  formula  of  the  acid  must  take  account 
of  this  fact.  Three  possibilities  present  themselves: 

,0-H  O  0  H         .O 


N 


0-H        H-C-O-H  H-C-O-Na       "&'     XO 


In  the  first  and  last  the  hydrogen  units  should  behave  alike.  The 
second  formula  is  the  only  one  which  expresses  the  replaceability  of 
one  unit  and  not  of  the  other  by  a  metal.  Since  the  hydrogen  in 
methane  is  not  replaceable  by  metals  (p.  588),  we  infer  that  the  unit 
directly  combined  with  carbon  is  the  non-replaceable  one.  Sodium 
formate  is  therefore  represented  by  the  third  formula. 

Acetic  Acid  HCOzCH3.  —  This  acid  is  produced  in  the  dry 
distillation  of  wood  (p.  610).  Large  quantities  of  it  are  manufactured 
from  dilute  alcohol  (p.  609).  The  properties  of  acetic  acid  have 
already  been  described  (p.  609). 

Although  four  atoms  of  hydrogen  are  contained  in  its  molecule, 
but  one  of  these  is  replaceable  by  metals.     This  fact  is  recognized  in 
the  constitutional  formula  (p.  563)  of  the  acid,  CH3CO(OH).     In  this 
acid  a  radical,  methyl  CH3—  ,  takes  the  place  of  H  in  formic  acid. 
O 
II 
Thus  the  group  —  C  —  O  —  H,  called  carboxyl,  is  contained  in  most 

carbon  acids,  and  in  each  of  them,  as  in  formic  acid,  bears  the 
replaceable  hydrogen  unit.  The  other  three  hydrogen  units  in  acetic 
acid,  however,  are  replaceable  by  chlorine,  as  is  the  case  with  the 
hydrogen  units  in  hydrocarbons. 

The  above  brief  statements  in  regard  to  the  mode  of  expressing  the  'chemical 
properties  of  a  substance  by  an  elaborated  formula  bring  out  a  tendency  which 
prevails  in  the  behavior  of  organic  substances  and  is  almost  entirely  lacking  in 
inorganic  chemistry.  The  units  may  be  removed  from  the  molecule  of  an  organic 
substance  one  by  one,  and  other  units  or  groups  may  be  substituted  for  them  with- 


616  INORGANIC  CHEMISTRY 

out  disturbing  the  rest  of  the  molecule.  The  changes  take  place,  not,  as  in  the  case 
of  ionized  substances,  by  the  splitting  of  the  molecule  into  two  or  more  groups 
which  act  as  wholes,  but  by  the  displacement  of  the  units  piecemeal  and  the  intro- 
duction of  new  properties  according  to  the  nature  of  the  groups  introduced.  Thus, 
if  by  any  means  we  replace  an  atom  of  hydrogen  by  the  carboxyl  radical  (see  above), 
the  product  is  an  acid.  If  we  replace  it  simply  by  the  group  OH  the  product  is 
an  alcohol.  Each  substitution  may  take  place  repeatedly  in  a  given  molecule,  so 
that  di-basic  or  tri-basic  acids,  di-hydric  or  tri-hydric  alcohols  (see  Glycerine),  or 
substances  which  contain  both  OH  and  —  COOH  in  the  same  molecule  (like 
lactic  acid  and  tartaric  acid),  are  formed.  Other  groups  which  may  be  introduced 
or  removed  are  —  NH2,  —  NO2,  —  CN,  etc.,  each  of  which  confers  upon  a  sub- 
stance the  properties  which  go  with  the  group,  irrespective  of  the  other  features 
which  the  structure  of  the  substance  may  already  present. 

Oxalic  Acid  flT2C2O4.  —  This  acid  is  dibasic,  and  consists  of  two 
carboxyl  groups.  Its  calcium  salt  is  the  least  soluble  of  the  salts  of 
calcium,  and  is  found  in  many  plants  in  the  form  of  bundles  of  needle- 
shaped  crystals.  Potassium-hydrogen  oxalate  KHC204  is  found  in 
the  juices  of  various  species  of  oxalis.  The  acid  may  be  made  by 
oxidation  of  sugar  with  nitric  acid.  The  white  crystals  used  in  the 
laboratory  are  the  hydrate  H2C204,2H20.  When  heated  carefully  it 
sublimes  unchanged.  Stronger  heating  decomposes  it  into  carbon 
dioxide  and  formic  acid,  and  the  latter  breaks  up,  in  part,  into  water 
and  carbon  monoxide.  In  the  presence  of  dehydrating  agents  like 
sulphuric  acid,  water  and  the  two  oxides  of  carbon  alone  are  formed 
(p.  577). 

ALCOHOLS,  ESTERS,  FATS,  SOAP,  AND  ETHERS 

Alcohols.  —  We  have  already  seen  that  when  wood  is  distilled, 

methyl  alcohol  is  found  in  the  fluid  product.     When  purified  this  is  a 

w  colorless  liquid  boiling  at  66°.     Its  solution  in 

T"  water    shows    no    evidence   of    ionization.     The 

rr  _  p  _  Q  _  TT    formula  (CH3OH)  makes  it  impossible  to  repre- 

•  sent  the  structure   of    the    substance    in    more 

TT  than  one  way.     All  alcohols  contain  the  group 

=  C  -  O  -  H  (cf.  p.  616). 

Common  alcohol,  ethyl  alcohol  C2H5OH  (p.  607),  is  a  member  of 
the  series  CnH2n+iOH.  There  are  also  many  alcohols  with  more  than 
one  OH  group  in  each  molecule.  Of  these,  the  one  we  shall  presently 
encounter  is  glycerine  C3H5(OH)3.  The  sugars  and  cellulose  are 
alcohols  with  several  hydroxyl  radicals. 


ALCOHOLS,  ESTERS,  FATS,   SOAP  617 

Esters.  —  When  an  organic  acid  and  an  alcohol  are  mixed,  a 
very  slow  chemical  action  takes  place,  which,  being  reversible,  in  no 
case  reaches  completion.  With  the  simplest  members  of  these  groups, 
formic  acid  and  methyl  alcohol,  for  example,  the  change  is : 

HCOOH  +  HOCH3  +±  HCOOCHs  +  H20. 

The  product  is  known  as  methyl  formate.  This  action  has  the 
appearance  of  a  neutralization  (p.  389),  but  is  different  in  several 
ways.  Alcohol  is  not  a  base,  and  in  aqueous  solution  it  does  not 
conduct  electricity.  Then,  true  neutralization  takes  place  instantly, 
while  the  foregoing  action,  and  all  like  it,  proceed  very  slowly.  Thus, 
although  acetic  acid  is  a  true  acid,  it  is  not  here  interacting  with  a 
base. 

The  corresponding  action  between  acetic  acid  and  ethyl  alcohol: 
CH3COOH  +  HOC2H5  ^  CH3COOC2H5  +  H2O,  results  in  the  forma- 
tion of  ethyl  acetate.  In  this  case,  when  equivalent  quantities  of 
the  initial  substances  have  been  used  without  any  solvent,  and  a 
condition  of  equilibrium  has  been  reached,  two-thirds  of  the  material 
is  found  to  have  been  transformed  into  ethyl  acetate  and  water. 
If  we  start  with  the  latter  materials  in  pure  form,  the  same  equilibrium 
point  is  reached,  and  one-third  of  the  material  is  converted  into 
acetic  acid  and  alcohol. 

The  products  are  named  as  if  they  were  salts,  and  are  sometimes 
called  ethereal  salts,  because  they  result  from  the  displacement  of 
the  hydrogen  of  an  acid  by  a  radical.  This  designation,  however, 
is  not  very  happy,  since  the  products  are  not  ionized  and  possess 
none  of  the  properties  of  salts.  The  special  name  esters,  therefore, 
has  been  given  to  them.  The  action  is  always  extremely  slow  and 
never  complete,  but  it  may  be  hastened  and  carried  to  completion  by 
the  introduction  of  some  substance  capable  of  absorbing  the  water 
and  so  preventing  the  reversal.  Concentrated  sulphuric  acid,  for 
example,  or  anhydrous  cupric  sulphate,  may  be  used. 

Inorganic  acids  also  interact  with  alcohols  to  give  esters.  Thus, 
nitroglycerine  (p.  528)  is  an  ester,  and  should  be  called  glyceryl  trini- 
trate.  The  use  of  sulphuric  acid  to  assist  in  the  removal  of  the  water 
is  illustrated  in  the  preparation  of  this  substance.  Guncotton 
(p.  528)  is  an  ester  of  nitric  acid  also,  for  cellulose  is  a  complex  alcohol. 
Ethyl-hydrogen  sulphate  (p.  591)  is  an  ester  of  sulphuric  acid.  In 
this  case  the  action  may  be  made  complete  by  using  sulphuric  acid 
containing  an  amount  of  sulphur  trioxide  sufficient  to  combine  with 
the  water  to  be  produced. 


618  INORGANIC   CHEMISTRY 

With  the  assistance  of  a  dehydrating  agent,  similar  actions  take 
place  between  any  alcohol  and  any  acid  (organic  or  inorganic). 
For  example: 

C3H5(OH)3  +  3H(C02CH3)  <=±  C3H5(CO2CH3)3  +  3H2O, 
glycerine  acetic  acid  glyceryl  acetate 

C3H5(OH)3  +  3H(C02C17H35)  <=±  C3H5(CO2C17H35)3  +  3H20. 
glycerine  stearic  acid  glyceryl  stearate 

The  glyceryl  radical  C3H5m  is  trivalent,  and  takes  the  place  of 
three  atoms  of  hydrogen. 

The  above  actions,  in  which  an  ester,  like  ethyl  acetate,  is  formed, 
may  be  almost  completely  reversed  if  a  sufficient  amount  of  water  is 
added  (cf.  p.  617).  The  hydrolysis  of  the  ester  is  hastened  by  the 
presence  of  free  acids  in  the  water.  This  is  owing  to  the  catalytic 
action  of  the  hydrogen  ions,  and  the  acceleration  is  proportional  to  the 
activity  of  the  acid  used.  The  acid,  however,  although  it  hastens  the 
action,  does  not  carry  it  beyond  the  condition  of  equilibrium  which  it 
would  eventually  have  reached  with  the  same  amounts  of  the  ester 
and  of  water  alone. 

When  an  ester  is  boiled  with  a  strong  base,  such  as  sodium 
hydroxide  solution,  the  salt  of  the  acid  and  an  alcohol  are  formed: 

CH3COOC2H5  +  NaOH  -»  CH3COONa  +  HOC2H5. 

With  more  complex  esters  the  sodium  salts  of  the  acids  thus  produced 
are  known  as  soaps,  and  this  general  kind  of  action  is  called,  therefore, 
saponification  (Lat.  sapo,  soap).  The  speed  with  which  it  proceeds 
may  be  used  as  a  means  of  measuring  the  activity  of  bases. 

Fats  and  Animal  and  Vegetable  Oils.  —  The  fats  and  oils 
found  in  animal  tissue,  or  pressed  from  seeds  of  plants,  are  com- 
posed mainly  of  esters.  Beef  fat  is  a  mixture  of  about  three-fourths 
glyceryl  palmitate  (palmitin)  C3H5(CO2Ci5H3i)3  and  glyceryl  stearate 
(stearin)  C3H5(CO2Ci7H35)3,  along  with  one-fourth  glyceryl  oleate 
(olein)  C3H5(CO2Ci7H33)3.  Lard  (hog  fat)  contains  a  much  larger 
proportion  of  the  last  (60  per  cent)  and  is  therefore  softer.  Butter 
contains  the  same  esters,  along  with  some  water  and  some  glyceryl 
butyrate  (butyrin)  C3H5(CO2C3H7)3.  Olive  oil  contains  much  olein 
(75  per  cent).  Cottonseed  oil  is  similar  in  composition,  and  is  used 
as  a  substitute  for  olive  oil,  or  for  butter  in  cooking. 

All  these  fats  and  oils  contain  certain  proportions  of  the  free 


ALCOHOLS,  ESTERS,  FATS,  SOAP  619 

organic  acids  (see  p.  625).  These  oils  must  not  be  confused  with 
mineral  oils,  which  are  mixtures  of  hydrocarbons. 

As  regards  physical  properties,  these  oils  are  all  insoluble  in  water, 
and  the  heavier  ones  also  in  cold  alcohol.  They  dissolve  readily, 
however,  in  ether,  benzine,  carbon  disulphide,  and  carbon  tetra- 
chloride.  Hence,  benzine  is  used  in  dry-cleaning  clothing  made 
of  silk  or  wool.  The  two  last  solvents  are  used  in  extracting  vege- 
table oils. 

Chemical  Properties  of  Fats  and  Oils.  —  All  fats  and  oils 
when  boiled  with  water,  and  more  rapidly  when  heated  (200°)  with 
water  in  a  closed  vessel,  are  decomposed.  The  ester  is  hydrolyzed, 
and  the  actions  in  the  four  equations  last  given  (pp.  617,  618)  are 
reversed.  Thus,  with  stearin: 

C3H5(CO2C17H35)3  +  3H20  ->  C3H5(OH)3  +  3HC02Ci7H35, 
stearin  glycerine  stearic  acid 

and  when  the  mixture  is  cooled,  the  acid,  being  insoluble  in  water, 
forms  a  solid  cake,  while  the  glycerine  is  in  solution  in  the  water. 
If  a  mixture  like  beef  fat  is  heated  with  water  in  this  way,  a  mix- 
ture of  palmitic,  stearic,  and  oleic  acids  is  obtained.  The  oleic 
acid  (liquid)  is  pressed  out,  and  the  residue  is  mixed  with  paraffin 
to  make  candles.  The  glycerine  is  separated  from  the  water  and 
used  in  making  nitroglycerine  (glyceryl  nitrate,  an  ester)  and  in 
medicine. 

When  fat  is  mixed  with  hot  sodium  hydroxide  solution,  it  first 
forms  an  emulsion,  in  which  the  fat  is  disseminated  in  minute  droplets 
through  the  liquid.  This  is  a  result  of  surface  tension.  When 
the  emulsion  is  boiled,  the  fat  is  slowly  decomposed  into  sodium 
palmitate,  stearate,  andoleate, and  glycerine.  The  change  is  precisely 
similar  in  plan  to  the  action  on  ethyl  acetate  (p.  618). 

C17H35COO  -  C  =  H2  HOCH2 

I  I 

Ci7H35COO-C-H  +  3NaOH  -» 3C,vH36COONa  +  HOCH 

I  I 

Ci7H35COO  -  C  =  H2  HOCH2 

stearin  sodium  stearate        glycerine 

When  common  salt  is  added  to  the  solution  ("  sal  ting  out")>  the 
sodium  salts  of  the  three  acids  (the  soap)  are  coagulated  and  separated 
as  a  floating  layer,  which  solidifies  when  cold.  The  glycerine  is 
contained  in  the  salt  solution. 


620  INORGANIC  CHEMISTRY 

With  potassium  hydroxide,  the  potassium  salts  are  obtained, 
and  constitute  soft  soap. 

The  soaps  are  purified  by  redissolving  in  water  and  again  salting 
out.  Dyes  and  perfumes  are  often  added.  Floating  varieties  are 
made  by  beating  the  soap  before  it  solidifies,  and  so  introducing 
bubbles  of  air.  Fine  sand  or  pumice  is  added  to  make  scouring 
soaps.  Mixing  with  glycerine  or  sugar  gives  transparent  soap. 

Chemical  Properties  of  Soaps.  —  Since  the  soaps  are  soluble 
salts  of  sodium,  they  are  largely  ionized  in  solution  and  interact 
with  acids  by  double  decomposition: 

Na(C02C17H35)  +  HC1  ->  NaCl  +  H(CO2C17H35H. 

The  acids  are  precipitated.  They  also  enter  into  double  decomposi- 
tion with  other  salts.  Thus,  hard  waters,  containing  compounds  of 
calcium  and  magnesium  in  solution,  give  precipitates  of  the  corre- 
sponding salts  For  example: 

2Na(C02C17H35)  +  CaSO4  -»  Na2S04  +  Ca(CO2Ci7H35)2|. 

Hence,  with  hard  water,  much  soap  is  wasted"  in  precipitating  the 
"  hardness." 

These  acids,  not  being  soluble  in  water,  have  no  effect  upon 
litmus;  but  the  fact  that  they  are  acids  may  be  recognized  when  it  is 
found  that  they  are  converted  into  soluble  salts  by  soluble  bases: 

C17H35COOHT  +  NaOH<^H20  +  C17H35COONa. 


Drying  Oils.  —  The  oils,  commonly  used  as  "dryers"  for  mixing 
with  varnish  and  paint  and  in  making  linoleum,  such  as  linseed  oil, 
hemp  oil,  poppy  oil,  and  nut  oil,  contain  esters  of  acids  with  unsatu- 
rated  radicals.  One  of  the  constituents,  for  example,  is  the  glyceryl 
ester  of  linoleic  acid.  The  formula  of  this  acid  is  C]7H3iCOOH.  It 
contains  four  hydrogen  atoms  less  than  the  corresponding  saturated 
acid  (stearic  acid).  These  oils,  especially  after  having  been  recently 
heated,  alone  or  with  catalytic  agents  like  lead  oxide  and  manganese 
dioxide,  absorb  oxygen  rapidly  from  the  air,  and  become  solid. 
They  do  not  dry,  in  the  ordinary  sense,  by  evaporation. 

Ether.  —  When  two  molecules  of  an  alcohol  lose  one  molecule 
of  water,  an  ether  is  produced: 

2CH3OH  -»  (CH3)2O  +  H2O. 


COLLOIDS.     CLEANSING  POWER  OF  SOAP  621 

Thus,  methyl  alcohol  gives  methyl  ether,  and  ethyl  alcohol,  ethyl  or 
common  ether.  The  action  is  most  easily  carried  out  by  two  steps. 
In  making  common  ether,  ethyl  alcohol  acts  upon  sulphuric  acid, 
giving  ethyl-hydrogen  sulphate  (p.  591  );  and  the  latter,  when 
warmed  gently  with  excess  of  alcohol,  gives  ethyl  ether: 

C2H5HS04  +  C2H5OH  -> (C2H5)20 1  +  H2SO4. 

The  ether  escapes  as  vapor  and  is  condensed. 

Ethyl  ether  is  a  volatile  liquid  boiling  at  35.6°.  It  is  largely  used 
as  a  solvent  for  iodine,  fats,  and  other  substances  not  readily  soluble 
in  water,  and  as  an  anaesthetic. 

COLLOIDAL  SUSPENSION.    CLEANSING  POWER  OF  SOAP 

Colloidal  Suspension.  —  To  explain  the  cleansing  power  of 
soap,  it  is  necessary  to  learn  more  about  colloids,  for  soap  in  solu- 
tion is  essentially  colloidal. 

The  simplest  colloidal  suspensions  are  those  of  metals  like  gold 
and  platinum.  They  can  be  made  by  forming  an  electric  arc  be- 
tween the  points  of  two  wires,  while  the  points  are  immersed  in 
water.  Liquids  of  various  colors,  depending  on  the  degree  of 
dispersion  (fineness  of  the  particles)  of  the  metal,  are  thus  formed. 
Such  a  liquid,  (1)  leaves  no  deposit  on  filter  paper,  (2)  shows  no 
elevation  in  the  boiling-point  of  the  solvent  and,  (3)  no  depression 
in  the  freezing-point.  (4)  The  suspended  body  has  little  or  no 
tendency  to  diffuse  into  a  layer  of  the  pure  solvent.  In  conse- 
quence, if  the  colloidal  solution  is  placed  in  a  diffusion-shell,  which 
is  a  test-tube  shaped  tube  of  filter-paper  or  parchment,  immersed 
in  water,  none  of  the  colloid  escapes  through  the  pores  of  the  shell. 
Ordinary  solutes  escape  more  or  less  quickly,  according  to  their 
molecular  weight.  Hence,  a  diffusion  shell  can  be  used  to  separate 
a  mixture  of  colloidal  and  non-colloidal  material.  Thus,  salt,  if 
present  with  colloidal  starch,  or  sugar  if  present  with  colloidal  gold, 
can  be  removed  by  changing  the  water  round  the  shell  until  no  more 
salt  (or  sugar)  is  found  to  come  out.  This  process  is  called  dialysis, 
and  was  devised  by  Graham. 

(5)  The  most  striking  property  of  colloids  is  shown  by  the 
ultramicroscope.  In  a  perfectly  darkened  room,  a  converging 
beam  of  strong  light  is  sent  horizontally  through  the  liquid  (Fig.  139) 
and  the  place  where  the  light  is  focussed  is  viewed  from  above, 
through  a  microscope.  Under  such  circumstances,  a  true  solution 


622  INORGANIC   CHEMISTRY 

remains  perfectly  dark,  but  a  colloidal  suspension  shows  minute  points 
of  light,  first  studied  by  Tyndall.  Colloidal  gold,  solutions  of  soap, 
starch,  gelatine,  and  dyes,  and  many  other  liquids 
exhibit  the  phenomenon.  The  points  of  light,  due 
to  particles  which,  although  minute,  contain  many 
molecules,  show  also  a  trembling  or  vibrating 
movement,  first  noticed  by  a  botanist  Brown 
(1827)  and  called  the  Brownian  Movement.  The 
motion  is  due  to  collisions  of  the  moving  mole- 
FIG.  139  cules  of  the  solvent  with  the  suspended  particles 

of  the  colloid  and,  when  the  suspension  is  very 
fine  (highly  "disperse")?  the  particles  shoot  about  rapidly. 
Other  properties  of  colloidal  suspensions  are  discussed  below. 

Theory  of  Colloidal  Suspension  and  Coagulation.  —  When 
wires  from  a  battery  are  immersed  in  the  liquid,  the  particles  of  a 
colloid  are  found  to  move  slowly  either  with  or  against  the  positive 
current.  The  phenomenon  is  called  electrophoresis.  Apparently, 
the  colloidal  particles  are  aggregates  of  molecules  of  an  insoluble 
substance,  collected  round  one  or  more  ions.  The  particles,  although 
relatively  large,  move  almost  as  rapidly  as  in  ionic  migration  (p.  345). 
This  affords  an  explanation  of  the  fact  that  the  particles  remain 
suspended,  and  do  not  settle.  They  are  individually  so  small  that 
they  are  kept  in  motion  by  collisions  with  the  molecules  of  the  solvent. 
If  they  could  unite  into  large  aggregates  —  like  the  particles  of  a 
precipitate  —  they  would  separate  like  any  ordinary,  insoluble 
substance.  But,  having  like  electrical  charges,  they  repel  one 
another,  and  so  remain  separate  and  in  suspension. 

Now  those  colloids  which  have  distinct  electrical  charges  can 
be  coagulated  or  flocculated,  and  so  precipitated  in  the  liquid,  by 
adding  a  solution  of  an  ionized  substance.  Thus,  colloidal  gold 
and  other  metals  are  negative,  and  an  equivalent  amount  of  a  posi- 
tive ion,  usually  H+,  is  present  also.  When  a  salt  is  added,  the 
positive  ion  of  the  salt  attaches  itself  to  the  negative  colloidal  me- 
tallic particles,  neutral  bodies  result,  coagulation  can  now  occur, 
and  precipitation  follows.  Bivalent  ions  are  more  effective  than 
univalent  ones  (see  Arsenic  trisulphide) .  Conversely,  a  positive 
colloid,  like  ferric  hydroxide,  is  coagulated  by  the  negative  ion  of 
a  salt,  and  more  easily  the  higher  the  valence  of  the  negative  ion. 
Furthermore,  one  colloid  will  coagulate  another  of  opposite  charge. 
Thus,  metaphosphoric  acid  is  a  negative  colloid  when  in  solution, 


COLLOIDS.     CLEANSING  POWER  OF  SOAP  623 

while  ortho-  and  pyrophosphoric  acid  are  not  colloidal.  Albumin 
is  usually  a  positive  colloid.  Hence  (p.  560),  metaphosphoric  acid 
and  albumin  coagulate  and  precipitate  one  another,  while  the  other 
two  acids  have  no  action  on  albumin. 

Starch  and  gelatine  are  neutral  colloids,  and  are  not  easily 
coagulated. 

Soap  Solution  Colloidal:  Salting  Out.  —  Soap  solution,  un- 
der the  ultramicroscope,  is  seen  to  contain  suspended  particles. 
A  test  with  litmus  also  shows  that  the  soap  is  partly  hydrolyzed: 

Na(C02R)  +  H20  +±  H(CO2R)  +  NaOH. 

Being  a  salt  of  a  little  ionized  acid,  the  negative  ion  of  the  soap 
tends  to  combine  with  the  H+  of  the  water:  H+ +  (CO2R)~  *± 
H(C02R),  leaving  the  ions  of  sodium  hydroxide.  Now  the  acid 
thus  set  free  combines  with  the  undissociated  molecules  of  the  salt 
to  form  an  acid  salt  NaH(C02R)2.  This  acid  salt  is  insoluble,  but 
remains  in  colloidal  suspension  as  a  negative  colloid.  When  a 
strong  solution  of  common  salt  (or  even  excess  of  sodium  hydrox- 
ide) is  added,  the  positive  ion  Na+  is  adsorbed  by  the  negative 
colloid  (the  acid  salt)  and  the  latter  is  coagulated.  In  coming  out 
as  a  precipitate,  it  seems  also  to  adsorb  most  of  the  sodium  hydrox- 
ide, so  that  the  precipitate  has  the  composition  of  soap. 

It  should  be  noted  that  salt  solution,  or  even  sodium  hydroxide 
solution,  will  coagulate  (salt  out)  the  soap  from  a  0.5  per  cent  solution, 
as  well  as  from  a  20  per  cent  or  a  stronger  solution.  This  is,  there- 
fore, not  a  case  of  precipitation  by  adding  excess  of  one  ion,  for  that 
occurs  only  in  saturated  solutions  (see  Chap.  XXXIV). 

The  Cleansing  Power  of  Soap.  Emulsions.  —  As  a  cleanser, 
soap  solution  —  or  suspension,  as  we  should  now  call  it  —  has  two 
properties.  It  removes  oil  and  grease  (insoluble  liquids)  by  forming 
an  emulsion  with  them.  It  also  removes  minute  solid  particles  of 
dirt,  by  taking  the  dirt  into  suspension  (next  section). 

When  an  oil,  such  as  kerosene,  is  violently  shaken  with  water, 
both  liquids  are  broken  into  minute  droplets,  and  an  opaque  mix- 
ture results.  The  droplets  of  each  liquid,  however,  quickly  join 
together  and  soon  the  mass  clears  up  and  shows  the  two  liquids  in 
separate  layers.  If,  however,  a  colloidal  suspension  is  used,  instead 
of  pure  water,  the  droplets  unite  much  more  slowly,  if  at  all,  and  a 
more  or  less  permanent,  opaque,  rather  viscous  mass  results.  Such 


624  INORGANIC  CHEMISTRY 

a  mixture  of  two  mutually  insoluble  liquids  is  called  an  emulsion. 
Thus,  a  few  drops  of  soap  solution  will  cause  the  kerosene  and  water 
to  remain  much  longer  in  the  condition  of  an  emulsion.  Similarly, 
vinegar  and  olive  oil,  when  vigorously  beaten  (French  dressing) 
separate  rather  quickly  into  two  layers.  But  if  the  yolk  of  an  egg 
(colloidal)  is  added  to  the  vinegar,  a  stiff,  almost  solid  mass  can  be 
made  (Mayonnaise  dressing),  which  will  remain  permanently  emulsi- 
fied. In  removing  grease,  therefore,  rubbing  with  soap  solution 
turns  the  grease  into  suspended  droplets  (emulsifies  it),  and  so  the 
grease  can  be  washed  away. 

This  behavior  of  a  colloid  can  be  explained.  When  the  kerosene 
and  water  are  divided  into  droplets,  with  a  great  increase  in  the 
total  surface,  and  in  the  surface  energy  of  both,  the  surface  tension 
of  water,  which  is  great,  favors  the  reunion  of  the  drops,  with 
diminishing  surface,  and  dissipation  of  the  surface  energy.  Now, 
while  ordinary,  dilute  solutions  have  a  surface  tension  close  to  that 
of  water,  colloidal  solutions  (such  as  0.5  per  cent  soap  solution) 
have  a  very  low  surface  tension.  Hence,  the  tendency  to  diminish 
the  surface  of  droplets  of  soap  solution,  by  coalescence,  is  slight 
and  ineffective.  Furthermore,  as  predicted  by  Willard  Gibbs 
of  Yale  University,  and  proved  by  experiment,  a  colloid  has  the 
peculiarity  that  it  tends  to  reach  a  higher  concentration  in  the 
surface  layer  than  it  has  elsewhere  in  the  liquid.  When  the  colloid 
has  adjusted  itself  to  a  state  of  equilibrium,  in  this  regard,  it  resists 
a  decrease  in  the  surface  (which  would  increase  its  concentration 
beyond  the  equilibrium  value),  just  as  much  as  it  resists  an  increase, 
which  would  diminish  its  concentration.  The  emulsion  of  a  colloidal 
suspension  with  an  immiscible  liquid  is  thus  a  stable  condition. 

Experiments  confirming  this  view  are  easily  made.  If  a  solu- 
tion of  a  dye  like  methyl  violet  (colloid)  be  shaken  violently,  and 
the  froth  (large  surface  in  proportion  to  quantity  of  liquid)  be  sepa- 
rated, it  is  found  that  the  liquid  produced  by  the  subsidence  of  the 
froth  (an  emulsion  with  air  is  not  permanent),  is  darker  in  shade, 
and  contains  more  dye,  than  an  equal  amount  of  the  original  liquid. 
Soap  solution,  after  being  shaken  likewise,  contains  relatively  more 
soap  in  the  froth  than  in  the  liquid. 

Adsorption  of  Colloidal  Matter.  —  As  we  have  seen  (p.  611), 
when  liquids  containing  colloidal  substances,  such  as  dyes  and 
natural  organic  coloring  matters,  are  shaken  with  pulverized  charcoal, 
the  colloid  is  adsorbed  by  the  charcoal  —  that  is,  it  adheres  to  the 


CYANOGEN  625 

surface  of  the  particles  of  the  charcoal.  This  principle  is  used  in 
decolorizing  sugar  (p.  606)  and  in  "bleaching"  oils.  Now,  soap  is 
also  removed  from  solution  (suspension)  by  shaking  with  charcoal  or 
with  infusorial  earth,  in  the  same  fashion. 

Pulverized  charcoal  is,  relatively,  a  coarse  powder.  If  soot, 
which  is  very  finely  divided  carbon,  be  freed  from  oil  or  grease  by 
washing  with  ether,  it  gives  a  loose,  non-caking  powder.  If  this 
powder  be  shaken  with  water,  it  settles.  If  it  be  shaken  with  dilute 
soap  solution,  it  remains  in  suspension,  and  the  liquid  resembles 
ink.  The  particles  are  so  fine  that,  instead  of  carrying  down  the 
colloidal  soap,  and  forming  a  precipitate,  as  charcoal  does,  they 
attach  themselves  to  the  colloidal  soap,  and  remain  suspended. 
This  is  therefore  adsorption,  with  the  difference  from  the  ordinary 
phenomenon,  that  the  colloid  carries  off  the  adsorbant,  instead  of  the 
adsorbant  carrying  down  the  colloid.  Now  dirt  is  composed  largely 
of  soot,  and  equally  fine  particles  of  other  substances.  Hence,  the 
soap  first  emulsifies  the  oil  on  the  hands  (or  on  soiled  linen)  and 
then  adsorbs  the  particles  of  dirt  which  are  thus  set  free. 

Formerly,  soap  solution  was  supposed  to  remove  grease  (and 
soot?)  because  of  its  slight  alkaline  reaction,  due  to  hydrolysis. 
This  explanation  must  be  given  up,  because:  (1)  an  alkali  so  dilute 
that  it  exists  in  equilibrium  with  the  free  fatty  acid,  can  not  possibly 
saponify  the  ester  contained  in  a  grease  spot.  (2)  Pure  alkali  of 
the  same  concentration  (or  stronger)  has  no  more  emulsifying  power 
than  has  water.  Such  an  alkaline  solution  will  indeed  emulsify  an 
animal  or  vegetable  oil  (cod-liver  oil,  cottonseed  oil,  castor  oil),  but  it 
does  so  by  interacting  with  the  free  fatty  acid  always  present  in 
such  oil  (p.  618)  and  forming  therefrom  a  soap.  Such  an  alkaline 
solution  does  not  emulsify  kerosene,  and  does  not  emulsify  natural 
oils  from  which  the  free  fatty  acids  have  been  removed  with  sodium 
hydroxide  solution,  although  soap  solution  does.  The  emulsifying 
agency  in  this  case  is  a  soap.  (3)  Very  dilute  alkali  has  no  more 
effect  upon  soot  than  has  water,  —  but  soap  solution  takes  clean 
(greaseless)  soot  instantly  into  permenent  suspension.  (4)  An 
aqueous  solution  of  saponin  C32H54Oi8,  obtainable  from  several 
plants,  although  it  contains  no  alkali,  lathers,  emulsifies,  and  adsorbs 
dirt,  just  as  does  soap.  It  is  a  colloid. 

CYANOGEN 

Cyanogen  C2-/V2.  —  This  compound,  being  endothermal,  is 
formed  in  small  amount  when  a  discharge  of  electricity  takes  place 


626  INORGANIC  CHEMISTRY 

between  carbon  poles  in  an  atmosphere  of  nitrogen  (cf.  p.  592). 
Cyanogen,  as  an  endothermal  substance,  is  more  easily  made  as  one 
product  in  an  action  which  as  a  whole  is  exothermal  (p.  312).  It 
is  prepared  by  allowing  a  solution  of  cupric  sulphate  to  trickle  into  a 
warm  solution  of  potassium  cyanide.  The  cupric  cyanide,  at  first 
precipitated,  quickly  decomposes,  giving  cuprous  cyanide  and 
cyanogen: 

2KNC  +  CuS04  -> Cu(NC)2  i  +  K2S04, 
2Cu(NC)2  -»  2CuNC  +  C2N2t . 

Cyanogen  is  a  very  poisonous  gas  with  a  characteristic,  faint  odor. 

Hydrocyanic  Acid  HNC.  —  This  acid,  called  also  prussic  acid, 

has  the  formula  HN  =  C,  and  is  most  easily  made  by  the  action  of  an 
acid  upon  a  cyanide  (see  Potassium  cyanide),  followed  by  distillation. 
It  is  a  colorless  liquid  boiling  at  26.5°.  It  has  an  odor  like  that  of 
bitter  almonds,  and  is  highly  poisonous.  In  aqueous  solution  it  is  an 
extremely  feeble  acid,  and  is  hardly  ionized  at  all.  In  consequence  of 
this,  potassium  cyanide  is  markedly  hydrolyzed  by  water,  and  its 
aqueous  solution  is  strongly  alkaline.  The  behavior  of  hydrocyanic 
acid  shows  it  to  be  an  unsaturated  body,  a  fact  which  is  taken  account 
of  in  the  above  formula,  and  illustrated  in  the  two  following  para- 
graphs. 

Cyanates  and  Thiocyanates.  —  When  potassium  cyanide  is 
fused  and  stirred  with  an  easily  reducible  oxide,  like  lead  oxide 
PbO,  the  metal  (for  example,  the  lead)  collects  at  the  bottom  of 
the  iron  crucible  in  molten  form,  and  potassium  cyanate  KNCO  is 
produced: 

KNC  +  PbO  -»  KNCO  +  Pb. 

Cyanic  acid  H  — N  =  C  =  0  is  very  unstable.  Ammonium  cya- 
nate NH4CNO  is  chiefly  remarkable  for  its  transformation  into  urea 
(p.  583). 

When  potassium  cyanide  in  aqueous  solution  is  boiled  with 
sulphur  or  with  a  polysulphide  (p.  421),  it  is  converted  into  potas- 
sium thiocyanate  KNCS.  This  salt,  or  ammonium  thiocyanate 
NH4NCS,  is  used  in  testing  for  ferric-ion  on  account  of  the  deep- 
red  color  of  ferric  thiocyanate  (cf.  p.  292).  The  ammonium  salt 
undergoes  at  170°  a  transformation  parallel  to  that  of  ammonium 
cyanate  (p.  583),  thiocarbamide  (sulpho-urea)  being  formed. 


FOODS 


627 


Fulminic  acid  H  — O  — N  =  C  (p.  541)  is  an  isomer  of  cyanic  acid 
(see  also  Calcium  cyanamide). 


FOODS 

Plants  and  animals  contain  substances  which  are  similar  in 
composition,  such  as  sucrose  and  lactose  (p.  605),  starch  and  glycogen 
(CeHioCWz,  animal  fats  and  vegetable  oils  (both  esters).  Albumins 
and  other  proteins  are  found  in  both.  They  differ,  however,  markedly 
in  the  sources  of  these  substances.  The  plant  uses  simple  materials, 
like  carbon  dioxide,  water,  and  potassium  nitrate.  The  animal  can 
make  no  use  of  these  substances  —  it  must  be  fed  on  complex  com- 
pounds. 

Foods.  —  Since  the  animal  is  continuously  eliminating  carbon 
dioxide,  moisture,  compounds  of  nitrogen,  salts,  and  other  sub- 
stances, and  is  also  giving  off  heat,  these  materials  must  be  replaced, 
and  fuel  must  be  furnished.  Like  the  plant,  an  animal  can  absorb 
only  dissolved  material.  But  it  prepares  its  own  solutions  in  a 
remarkable  laboratory,  the  digestive  tract.  The  production  of  suitable 
soluble  substances  is  called  digestion. 

The  following  table  shows  the  chief  components  of  animal  food, 
and  the  proportions  in  which  they  are  present  in  the  chief  foods 
used  by  man: 


Water. 

Protein. 

Fat. 

Carbo- 
hydrate. 

Ash. 

Beef  (lean) 

73  8 

22   1 

2  9 

1  2 

Cod 

82  6 

15  8 

0  4 

1.2 

73.7 

14  8 

10  5 

1.0 

Mi  Ik* 

87.0 

3.3 

4.0 

5.0 

0.7 

Butter           

11.0 

1.0 

85.0 

3.0 

Cheese  (cheddar)    

27.4 

27.7 

36.8 

4.1 

4.0 

Oatmeal     

7.3 

16.1 

7.2 

67.5 

1.9 

Wheat  flour 

11  9 

13  3 

1  5 

72  7 

0.6 

Beans  (dried)                               .   .    . 

12  6 

22  5 

1.8 

59.6 

3.5 

Almonds                   

4.8 

21.0 

54.9 

17.3 

2.0 

Maize  (green  corn)     

75.4 

3.1 

1.1 

19.7 

0.7 

Potatoes    

78.3 

2.2 

0.1 

18.4 

1.0 

Lettuce 

94  7 

1  2 

0  3 

2.9 

0.9 

*  The  emulsified  fat  separates  slowly  as  the  cream;  the  protein  (casein,  colloidally  suspended  in 
the  skim  milk)  is  coagulated  by  rennet  and  constitutes  cheese;  the  carbohydrate  (lactose,  a  sugar) 
w  then  left  in  the  water,  along  with  inorganic  salts. 


628  INORGANIC  CHEMISTRY 

We  observe  that  the  common  animal  foods,  except  milk,  containing 
lactose  (p.  605),  carry  no  carbohydrates  (the  ox  liver  contains  about 
2  per  cent  of  glycogen);  that  potatoes  and  corn,  when  dried,  are 
nearly  all  starch;  that  lean  meat,  dry,  is  all  protein;  that  some 
seeds  (wheat  and  beans)  contain  little  fat,  some  (oats)  much  more 
fat,  and  some  (almonds  and  nuts)  a  large  amount;  and  that  lettuce 
is  mainly  water,  with  useful  inorganic  salts  in  solution,  and  cellulose. 
The  proteins,  several  of  which  have  been  mentioned  (pp.  520, 
528,  547)  are  white,  amorphous  substances  containing,  besides 
carbon,  hydrogen,  and  oxygen,  a  large  proportion  of  nitrogen  (16 
per  cent),  some  sulphur  (1  per  cent)  and  frequently  iron  and  phos- 
phorus as  well. 

Digestion.  —  The  process  of  rendering  the  constituents  of  food 
soluble  is  like  fermentation  (p.  607)  —  it  is  performed  mainly  by 
enzymes.  Each  class  of  components  is  handled  by  one  or  more 
enzymes.  Thus,  starch  (in  bread  and  potatoes)  is  partly  digested 
during  mastication  by  ptyalin  (an  amylase,  p.  608)  in  the  saliva, 
and  partly  by  amylopsin  in  the  small  intestine.  The  resulting 
maltose  is  decomposed  into  glucose  by  another  enzyme,  and  passes 
into  the  blood,  where  it  is  oxidized,  furnishing  heat.  In  diabetes, 
much  of  the  glucose  escapes  oxidation,  and  is  eliminated.  Again, 
the  fats  are  hydrolyzed  into  the  acids  and  glycerine  by  lipases 
(fat-splitting  enzymes)  in  the  bile,  and  the  acids  go  into  solution 
(probably  colloidal).  The  acids  and  glycerine  recombine  to  form 
fats  in  the  blood,  and  are  either  deposited  in  the  tissues  or  oxidized. 
Finally,  the  proteins  are  changed  in  a  similar  way  into  peptones 
which  are  soluble  in  water,  and  in  this  form  are  able  to  pass  through 
the  wall  of  the  intestine. 

Fuel  Value  of  Food.  —  Although  food  is  required  to  replace 
waste,  much  of  it  is  needed  to  furnish  energy,  by  its  oxidation,  so 
that  muscular  movements  may  be  maintained,  and  the  tempera- 
ture of  the  body  kept  up  to  its  normal  value  (37°  C.).  Thus,  the 
fuel  values  of  foods  are  important.  The  average  fuel  values,  ex- 
pressed in  large  calories  (1  Cal.  =  1000  cal.  as  previously  defined, 
p.  98),  per  gram,  are: 

Carbohydrates,  4  Cal.     Fats,  9  Cal.     Proteins,  4  Cal. 

The  fuel  values  per  pound  (453.6  g.)  are  453.6  times  greater. 

Healthy  life  cannot  be  maintained  on  one  kind  of  food  —  a 


FOODS  629 

mixed  diet  is  necessary.  In  general,  it  is  held  that  100  g.  of  pro- 
teins (giving  400  Cal.)  per  day,  and  a  sufficient  amount  of  other 
foods  to  give  a  total  fuel  value  of  2200  Cal.  is  enough  for  a  person 
doing  no  physical  labor.  When  physical  labor  is  involved,  larger 
values,  up  to  3800  cal.  per  day,  are  necessary.  The  data  in  the 
table  (p.  627)  will  enable  one  to  calculate  the  fuel  value  of  100  g. 
(or  of  1  Ib.)  of  each  kind  of  food. 

Exercises.  —  1.  Make  the  graphic  formulae  of  methyl  acetate 
(p.  615),  ethyl  formate,  ethylene  bromide  (p.  591),  oxalic  acid 
(p.  616),  ethyl  ether  (p.  621). 

2.  Make  equations  for  the  hydrolysis  of  starch  to  maltose  (p.  608), 
the  saponification  of  olein  (p.  618). 

3.  Prepare  a  summary  of  the  various  statements  that  have  been 
made  in  the  text  about  catalysis  (e.g.,  pp.  97,  128,  217,  222,  413,  433, 
445,  516,  570,  606,  607,  618),  and  illustrate  fully. 

4.  Calculate  the  fuel  value  of  1  Ib.  each  of  (a)  oatmeal,  (6) 
potatoes,  (c)  lettuce. 

5.  Calculate  the  weights,  both  in  pounds  and  in  grams,  of  100 
Cal.  portions  of,  (a)  eggs,  (6)  wheat  flour,  (c)  almonds,  (d)  lettuce. 

6.  At  current  market  prices,  what  would  be  the  cost  per  100  Cal. 
portion  of  beef,  cod,  butter,  and  wheat  flour,  respectively? 

7.  Why  are  there  no  colloidal  suspensions  of  iron  or  zinc  in 
water? 


CHAPTER  XXXI 
SILICON  AND  BORON 

IN  respect  to  chemical  relations  there  is  a  close  resemblance 
between  silicon  and  carbon.  Silicon  gives  a  monoxide,  but  is  quadri- 
valent in  all  its  other  compounds.  It  is  strictly  a  non-metallic 
element. 

Occurrence.  —  Silicon,  unlike  carbon,  is  not  found  in  the  free 
condition.  In  combination  it  is  the  most  plentiful  element  after 
oxygen,  and  constitutes  more  than  one-quarter  of  the  crust  of  the 
earth.  The  oxide  is  silica  or  sand  SiC>2,  and  this  oxide  and  its  com- 
pounds are  components  of  many  rocks.  In  the  inorganic  world 
silicon  is  the  characteristic  element  to  almost  as  great  an  extent  as  is 
carbon  in  the  organic  realm. 

Preparation.  —  When  finely  powdered  magnesium  and  sand  are 
mixed,  and  one  part  of  the  mass  is  heated,  a  violent  action  spreads 
rapidly  through  the  whole : 

2Mg  +  Si02  ->  Si  +  2MgO. 

At  the  same  time,  and  especially  if  excess  of  the  metal  is  used,  some 
magnesium  silicide  Mg2Si  is  formed  also.  The  mixture  is  treated 
with  a  dilute  acid  which  decomposes  the  magnesium  oxide  and  the 
silicide,  and  leaves  the  silicon  (amorphous)  undissolved. 

When  amorphous  silicon  is  dissolved  in  molten  zinc,  the  mass, 
after  solidification,  is  found  to  contain  crystalline  silicon.  This  form 
may  be  made  in  one  operation,  by  heating  three  parts  of  potassium 
fluosiHcate  K2SiF6  with  one  part  of  sodium  and  four  parts  of  zinc  in  a 
closed  crucible.  The  sodium  displaces  the  silicon  and  combines 
with  the  fluorine,  while  the  zinc  acts  as  a  solvent  as  before.  The 
zinc  is  removed  by  the  action  of  a  dilute  acid,  the  silicon  remaining 
unaffected. 

Silicon  and  ferrosilicon  (an  alloy  of  iron  and  silicon)  are  now 
made  on  a  large  scale,  the  former  by  heating  sand  and  carbon,  the 
latter  by  heating  a  mixture  of  ferric  oxide  and  sand  with  carbon 
in  the  electric  furnace  (p.  569) . 

630 


SILICON  AND  BORON  631 

Properties.  —  Amorphous  silicon  is  a  brown  powder.  It  unites 
with  fluorine  at  the  ordinary  temperature,  with  chlorine  at  430°, 
with  bromine  at  500°,  with  oxygen  at  400°,  with  sulphur  at  600°,  with 
nitrogen  at  about  1000°,  and  with  carbon  and  boron  at  temperatures 
attainable  only  in  the  electric  furnace.  It  is  slowly  oxidized  by 
aqua  regia  to  silicic  acid,  and  is  dissolved  by  a  mixture  of  hydro- 
fluoric acid  and  nitric  acid,  giving  silicon  tetrafluoride. 

Crystallized  silicon  consists  of  black  needles  belonging  to  the 
regular  system,  and  is  less  active  than  the  other  variety.  It  oxidizes 
superficially  at  400°,  and  the  dioxide  formed  on  the  surface  hinders 
further  action.  With  chlorine  and  fluorine  it  unites  easily  when 
heated.  Gaseous  hydrogen  fluoride  interacts  violently  with  it  at  a 
high  temperature,  heat  and  light  are  given  out,  and  silicon  tetra- 
fluoride and  hydrogen  are  formed.  It  is  slowly  attacked  by  hydro- 
fluoric acid  mixed  with  nitric  acid,  but  not  by  any  others  of  the 
oxygen  acids. 

Silicon  and  ferrosilicon  act  readily  upon  a  cold  solution  of  sodium 
hydroxide  (cf.  p.  122),  the  metasilicate  being  formed: 

Si  +  2NaOH  +  H20  -*  Na2Si03  +  2H2T- 

This  is  one  of  the  sources  of  hydrogen  for  filling  balloons  and  air- 
ships. A  layer  of  petroleum  on  the  surface  prevents  frothing. 
Hydrogenite  is  a  mixture  of  ferrosilicon  and  dry  sodium  hydroxide, 
to  which  water  is  added  drop  by  drop.  Three  kilograms  give  one 
cubic  meter  of  hydrogen. 

Silicon  seems  to  be  more  active  than  carbon,  for,  when  it  is 
heated  with  fused  potassium  carbonate,  potassium  silicate  is  formed 
and  carbon  is  liberated.  Both  kinds  of  silicon  differ  from  carbon  in 
being  fusible  at  a  very  high  temperature.  The  solidified  product 
is  crystallized  silicon. 

Silicon  Hydride  SiH^.  —  Silicon  differs  from  carbon  in  giving 
two  well-defined  compounds  with  hydrogen,  SiH4  and  Si2H6.  The 
former  is  liberated  as  a  gas  by  the  action  of  hydrochloric  acid  upon 
magnesium  silicide: 

Mg2Si  +  4HC1  ->  2MgCl2  +  SiH4  T- 

The  action  is  similar  to  that  by  which  hydrogen  sulphide  is  made. 
Since  the  magnesium  silicide  always  contains  free  magnesium,  hydro- 
gen is  liberated  at  the  same  time.  By  leading  the  gases  through  a 
tube  surrounded  by  liquid  air,  the  silicon  hydride  is  reduced  to  liquid 


632  INORGANIC  CHEMISTRY 

form,  while  the  hydrogen  passes  on.  The  mixture  with  hydrogen 
is  spontaneously  inflammable.  The  pure  gas  becomes  so  only  when 
its  pressure  is  reduced  (cf.  p.  550).  In  the  air,  however,  it  is  easily 
inflammable,  by  contact  with  a  warm  body.  When  heated,  it  de- 
composes into  its  constituents. 

Silicon  Tetrachloride  SiCl^.  —  This  compound  is  made  by 
direct  union  of  the  free  elements.  It  is  more  conveniently  pre- 
pared by  passing  chlorine  over  a  strongly  heated  mixture  of  silicon 
dioxide  and  carbon.  The  gaseous  products  enter  a  condenser  in 
which  the  tetrachloride  assumes  the  liquid  form: 

2C12  +  SiO2  +  2C  ->  SiCl4  +  2CO. 

Chlorine  is  unable  to  displace  oxygen  from  combination  with  silicon, 
and  has,  therefore,  when  alone,  no  effect  upon  sand.  In  the  above 
action,  therefore,  the  carbon  is  used  to  secure  the  oxygen  while  the 
chlorine  combines  with  the  silicon.  This  kind  of  interaction  is  in 
some  degree  different  from  any  which  we  have  hitherto  encountered. 
The  principle  underlying  it  was  formerly  used  for  making  many 
chlorides  (e.g.,  BC13,  A1C13,  CrCl3)  from  oxides,  before  simple  ways  of 
obtaining  the  elements  in  the  free  condition  were  known. 

Silicon  tetrachloride  is  a  colorless  liquid  (b.-p.  59°)  which  fumes 
strongly  in  moist  air,  giving  silicic  acid  Si(OH)4.  It  acts  violently 
upon  cold  water,  and  in  this  respect  differs  from  carbon  tetrachloride : 

SiCU  +  4H2O  ->  4HC1  -f  Si(OH)4 . 

The  silicic  acid  (q.v.)  soon  appears  as  a  gelatinous  precipitate. 

When  silicon  is  heated  in  a  stream  of  dry  hydrogen  chloride,  a 
mixture  of  silicon  tetrachloride  and  silico-chloroform  SiHCl3  is  pro- 
duced. The  latter  is  a  volatile  liquid  boiling  at  34°. 

Silicon  Tetrqfluoride  SiF4. —  When  strong  hydrofluoric  acid 
acts  upon  sand,  this  gas  is  liberated : 

Si02  +  4HF  ->  2H20  +  SiF4T. 

Since  the  water  interacts  with  the  tetrafluoride  (see  below),  the  latter 
is  usually  made  by  the  use  of  powdered  calcium  fluoride  and  excess 
of  sulphuric  acid.  In  this  way  the  hydrogen  fluoride.is  generated  in 
contact  with  the  sand,  and  at  the  same  time  the  sulphuric  acid  takes 
possession  of  the  water.  Hydrofluoric  acid  acts  in  a  corresponding 
way  upon  all  silicates  (q.v.),  whether  these  are  minerals  or  are  artificial 
silicates  like  glass  (cf.  p.  283). 


SILICON  AND  BORON  633 

Silicon  tetrafiuoride  is  a  gas  which  becomes  solid,  without  liquefy- 
ing (cf.  p.  574),  when  cooled  to  —102°.  It  fumes  strongly  in  moist 
air,  and  acts  vigorously  upon  water.  This  interaction  is  different 
from  that  of  the  tetrachloride,  because  the  excess  of  the  tetrafluoride 
forms  a  complex  compound  with  the  hydrofluoric  acid: 

SiF4  +  4H20  ->  Si(OH)4  (+  4HF)  (1) 

(4HF)  +  2SiF4  -»  2H2SiF6 •  (2) 

3SiF4  +  4H2O  ->  Si(OH)4  +  2H2SiF6 

The  silicic  acid  is  precipitated  in  the  water,  and  may  be  separated  by 
filtration,  leaving  a  solution  of  hydrofluosilicic  acid. 

Hydrofluosilicic  Acid  H2SiF6.  —  This  acid  is  stable  only  in  solu- 
tion. When  the  water  is  partly  removed  by  evaporation,  silicon 
tetrafluoride  is  given  off,  while  most  of  the  hydrogen  fluoride  remains 
to  the  last.  Its  salts  are  decomposed  in  a  corresponding  way  when 
they  are  heated.  This  acid  is  used  in  analysis,  chiefly  because  its 
potassium  salt  is  one  of  the  few  salts  of  this  metal  which  are  relatively 
insoluble  in  water.  The  barium  salt  is  also  insoluble,  but  most  of  the 
salts  of  the  heavy  metals  are  soluble. 

Silicon  Dioxide  SiO2.  —  This  substance  may  be  made  in  the 
form  of  a  white  powder  by  heating  precipitated  silicic  acid.  It  is 
found  in  many  different  forms  in  nature.  In  large,  transparent,  six- 
sided  prisms,  with  pyramidal  ends  (Fig.  2),  it  is  known  as  quartz  or 
rock  crystal.  When  colored  by  manganese  and  iron  it  is  called 
amethyst,  when  by  organic  matter,  smoky  quartz.  A  special  arrange- 
ment of  the  structure  gives  cat's  eye.  Amorphous  forms  of  the  same 
material,  often  colored  brown  or  red  with  ferric  oxide,  are  agate, 
jasper,  and  onyx,  the  last  much  used  in  making  cameos.  Infusorial 
or  diatomaceous  earth  (Tripoli  powder)  is  composed  of  the  tests  of 
minute  organisms,  and  is  used  in  scouring  materials  and  for  decolor- 
izing oils.  Sponges  are  also  made  of  silica  of  organic  origin.  Slightly 
hydra  ted  forms  of  silica  are  the  opal  and  flint.  A  crystalline  variety 
belonging  to  the  hexagonal  system,  but  showing  entirely  different 
crystalline  forms,  occurs  occasionally  in  minute  crystals,  and  is  called 
tridymite. 

Silica  is  found  in  the  hard  parts  of  straw,  of  some  species  of  horse- 
tail (equisetum) ,  and  of  bamboo.  In  the  form  of  whetstones  it  is  used 
for  grinding.  The  clear  crystals  are  employed  in  making  spectacles 
and  optical  instruments,  and  are  more  transparent  to  ultra-violet 


634  INORGANIC  CHEMISTRY 

light  than  is  glass.  Pure  sand  is  used  in  glass  manufacture  (q.v.). 
Infusorial  earth,  on  account  of  the  tubular  form  of  many  of  the 
minute  structures  of  which  it  is  composed,  can  absorb  three  times  its 
own  weight  of  nitroglycerine,  and  is  therefore  employed  in  making 
dynamite.  Recently,  pieces  of  chemical  apparatus  have  been  manu- 
factured by  fusing  quartz  (m.-p.  1600°)  in  the  oxy-hydrogen  flame, 
or  the  electric  furnace.  The  material  does  not  crystallize  on  cooling, 
and  is  amorphous,  like  glass.  Owing  to  the  low  coefficient  of  expan- 
sion of  silica,  the  vessels  can  be  heated  to  a  red  heat  and  chilled  in 
cold  water,  without  risk  of  fracture. 

Silicates.  Water  Glass.  —  Silicon  dioxide,  although  differing 
profoundly  from  carbon  dioxide  in  its  physical  nature,  nevertheless 
behaves  like  the  latter  chemically.  Thus,  when  boiled  with  sodium 
hydroxide  solution  it  forms  sodium  metasilicate  Na^SiOs  or  ortho- 
silicate  Na4SiO4. 

Si02  +  2NaOH  ->  Na^SiOs  +  H20. 


The  salt  is  left  as  a  gelatinous  solid  ("  soluble  glass")  when  the 
water  is  evaporated.  The  silicates  of  potassium  and  sodium  may 
also  be  obtained  by  boiling  sand  with  the  carbonates  of  these  metals, 
or,  more  rapidly,  usually  as  metasilicates  (see  below),  \yy  fusing  the 
mixture: 

Si02  +  K2C03  -rf  K2Si03  +  C02T. 

Water  glass  or  soluble  glass,  being  a  salt  of  a  feeble  acid  with  an 
active  base,  gives  an  alkaline  solution  (pp.  399,  559).  When  man- 
ufactured for  commercial  use,  it  has  the  composition  Na2Si205 
(Na2SiO3,SiO2),  which  is  less  alkaline.  It  is  used  as  a  filler  in  cheap 
soaps,  for  fireproofing  and  waterproofing  timber  and  textiles,  and  for 
preserving  eggs. 

Silicic  Acid  H^SiO^.  —  When  acids  are  added  to  a  solution  of 
sodium  silicate,  silicic  acid  is  set  free.  After  a  little  delay  it  usually 
appears  as  a  gelatinous  precipitate.  When,  however,  the  silicate  is 
poured  into  strong  hydrochloric  acid,  no  precipitation  occurs.  The 
silicic  acid  remains  in  colloidal  suspension: 

Na4Si04  +  4HC1  -»  4NaCl  +  Si(OH)4, 
Na2Si03  +  2HC1  +  H2O  ->  2NaCl  +  Si(OH)4. 

The  gelatinous  precipitate,  when  it  has  been  dried,  contains  a  smaller 
proportion  of  the  elements  of  water.  There  seem  to  be  no  definite 


SILICON  AND  BORON  635 

stages,  indicating  the  existence  of  various  acids,  such  as  we  observe 
with  phosphoric  acid.  We  should  expect  the  vapor  tension  of 
the  water  to  decrease  by  steps,  each  of  which  should  correspond  to 
some  acid  of  a  particular  degree  of  hydration  (cf.  p.  560),  but,  probably 
because  it  is  amorphous  while  the  phosphoric  acids  are  crystalline, 
nothing  of  the  sort  is  observed.  The  final  product  of  drying  is  the 
dioxide. 

Silicic  acid  is  a  very  feeble  acid.  For  this  reason  silicic  acid  can 
be  completely  displaced  from  its  salts,  even  by  so  weak  an  acid  as 
carbonic  acid.  Since  the  water  in  some  geyser  regions  contains 
alkali  silicates,  the  action  of  the  carbon  dioxide  in  the  air  causes 
deposition  of  silica  round  the  places  where  the  water  issues  from  the 
ground.  This  form  is  known  as  siliceous  sinter.  Striking  scenic 
effects,  as  in  the  white  and  pink  terraces  of  New  Zealand,  are  some- 
times produced  by  this  method  of  deposition. 

The  suspension  of  colloidal  silicic  acid  can  be  freed  from  the  acid 
and  sodium  chloride  (see  equation,  above)  by  dialysis  (p.  621). 
It  is  a  positive  or  a  negative  colloid,  according  to  the  mode  of  prep- 
aration, and  the  two  kinds  are  coagulated  by  addition  of  salts  having 
bivalent  negative  and  positive  ions,  respectively. 

Mineral  Silicates.  —  While,  in  the  absence  of  definite  knowl- 
edge, silicic  acid  is  presumed  to  be  the  ortho-acid  Si(OH)4,  and  no 
other  silicic  acids  have  been  made,  the  salts  are  most  easily  classified 
by  imagining  them  to  be  derived  from  various  acids  representing 
different  degrees  of  hydration  of  the  dioxide,  or,  to  put  it  the  other 
way,  different  degrees  of  dehydration  of  the  ortho-acid.  The  follow- 
ing equations  show  the  relation  of  the  ortho-acid  to  some  of  the 
silicic  acids  whose  salts  are  most  commonly  found  amongst  minerals : 

H4SiO4-    H2O-^H2Si03  (  =    H2O,SiO2)     Metasilicic  acid. 
2H4Si04  -    H20  -*  H6S207  ( =  3H20,2Si02)  j  Di  m  .       . , 
2H4Si04  -  3H20  ->  H2Si206  ( =    H2O,2SiO2)  i  ~ 
3H4SiO4  -  4H2O  ->  H4Si3O8  (=  2H2O,3SiO2)   Trisilicic  acid. 

Di-  and  trisilicates  are  those  derived  from  acids  containing  two  and 
three  units  of  silicic  anhydride,  respectively,  in  the  formula.  The 
valences  of  the  radicals  of  the  acids  are  shown  by  the  number  of 
hydrogen  units  in  the  formulae. 

The  composition  of  minerals  is  often  exceedingly  complex.  This  is 
due  to  the  fact  that  amongst  them  mixed  salts  (p.  401)  are  very 
common,  in  which  the  hydrogen  of  the  imaginary  acid  is  displaced  by 


636  INORGANIC  CHEMISTRY 

two  or  more  metals  in  such  a  way  that  the  total  quantity  of  the  metals 
is  equivalent  to  the  hydrogen.  The  following  list  presents  in  tabular 
form  some  typical  or  common  minerals,  arranged  according  to  the 
above  mentioned  classification: 

ORTHOSILICATES    (H4Si04)  METASILICATES    (H2Si03) 

Zircon,  ZrSiO4  Wollastonite,         CaSiO3 

Garnet,  CaaAl2(Si04)s  Beryl,  GlaAl^SiC^e 

Mica,  KH2Al3(SiO4)3  Talc  (soapstone),  H2Mgs(Si03)4 

Kaolin,  H2Al2(SiO4)2,H2O  Asbestos,  Mg3Ca(SiO3)4 

DISILICATE    (H6Si2O7)  TRISILICATE    (H4Sl308) 

Serpentine,  Mg3Si2O7,2H2O          Orthoclase  (felspar),  KAlSi3O8 

It  will  be  seen  that  the  total  valence  of  the  metal  units  is  equal  to 
that  of  the  acid  radicals.  Thus,  in  beryl  there  are  six  equivalents  of 
glucinum  (beryllium)  and  six  of  aluminium,  taking  the  place  of  twelve 
units  of  hydrogen  in  (H2SiO3)6. 

Garnets  are  pulverized  in  manganese-steel  crushers  and  used  in 
making  sandpaper.  Mica,  which  is  obtained  in  large  sheets  from 
Farther  India,  is  used  in  making  lamp-chimneys,  for  windows  in 
stoves,  and  as  an  insulator  in  electrical  apparatus.  Kaolin,  or  clay, 
like  mica,  is  an  acid  orthosilicate.  It  is  formed  in  nature  as  the 
result  of  the  action  of  water  and  carbonic  acid  upon  minerals  like 
felspar.  In  such  cases,  those  elements  which  can  form  carbonates 
like  potassium,  magnesium,  and  calcium,  are  usually  displaced  from 
combination  with  the  silicic  acid.  Aluminium,  however,  is  too 
feeble  a  base  to  form  a  carbonate,  and  is  thus  left  in  combination 
as  silicate.  A  clay  containing  lime  is  called  a  marl,  and  one  contain- 
ing sand,  a  loam. 

Some  of  these  minerals  frequently  occur  mixed  together  as  regular 
components  of  certain  igneous  rocks.  Thus,  granite  (p.  4)  is  a  more 
or  less  coarse  mixture  of  quartz,  mica,  and  felspar.  Frequently 
the  oblong,  flesh-colored  or  white  crystals  of  the  last  are  large  and 
very  conspicuous  both  in  granite  and  in  porphyry.  In  basalt  the 
components  are  usually  less  easily  visible  to  the  eye.  Lava  is  the 
name  for  any  rock  recently  ejected  from  a  volcano.  Pumice-stone 
is  the  name  given  to  the  parts  of  the  lava  which  are  porous,  having 
acquired  this  texture  from  the  expansion  of  bubbles  of  gas  consequent 
upon  release  of  pressure.  Sandstone  is  composed  of  sand  cemented 
together  by  clay  or  by  calcium  carbonate,  and  colored  brown  or  yellow 
by  ferric  oxide. 


SILICON  AND  BORON  637 

The  high  melting-point  of  silica,  compared  with  carbon  dioxide, 
and  the  formation  of  these  complex  silicates,  indicate  that  the  oxide 
is  highly  associated  (SiC^)*. 

The  hydrated  silicates  are  decomposed  by  hydrochloric  acid, 
but  the  others  are  not  affected  by  acids,  or  are  affected  with  extreme 
slowness.  Hence  special  means  have  to  be  taken  to  get  the  con- 
stituents of  such  minerals  into  soluble  form  for  the  purpose  of  analysis. 
Two  methods  are  in  use.  Sometimes  the  finely  powdered  mineral  is 
heated  in  a  platinum  dish  with  hydrofluoric  acid  until  all  the  silicon 
has  passed  off  in  the  form  of  the  tetrafluoride.  Since  the  use  of  the 
fluorides  would  lead  to  difficulties  in  the  course  of  the  analysis,  the 
resulting  mixture  of  fluorides  of  the  mentals  is  next  heated  strongly 
with  concentrated  sulphuric  acid,  and  the  mixture  of  sulphates  thus 
produced  is  treated  according  to  the  usual  routine.  In  other  cases  the 
finely  powdered  mineral  is  fused  at  a  bright  red  heat  with  a  mixture  of 
potassium  and  sodium  carbonates.  In  this  way  carbonates  of  the 
metals  are  formed,  along  with  potassium  and  sodium  silicate.  Treat- 
ment with  water  dissolves  the  latter,  and  leaves  the  carbonates  to  be 
handled  in  the  usual  way. 

BORON 

As  regards  chemical  relations,  boron,  being  a  uniformly  trivalent 
element,  is  a  member  of  the  aluminium  family.  Yet  it  is  a  pro- 
nounced non-metallic  element,  and  its  oxide  and  hydroxide  are 
almost  wholly  acidic;  aluminium  is  a  metal,  and  with  its  oxide  and 
hydroxide  basic  properties  predominate.  Boron  and  its  compounds 
really  resemble  carbon  and  silicon  and  their  compounds  in  all  chemical 
properties,  excepting  that  of  valence. 

Occurrence.  —  Like  silicon,  boron  is  found  in  oxygen  com- 
pounds, namely,  in  boric  acid  (q.v.)  and  its  salts.  Of  the  latter, 
sodium  tetraborate  Na^O?,  or  borax,  came  first  from  India  under 
the  name  of  tincal.  It  constitutes  a  large  deposit  in  Borax  Lake  in 
California.  Colemanite,  Ca2B6Oij,5H20,  from  California,  and  other 
complex  borates,  furnish  a  large  part  of  the  commercial  supply  of 
compounds  of  boron. 

Preparation.  —  When  boric  oxide  is  heated  with  powdered  mag- 
nesium: B2O3  +  3Mg  — >  3MgO  +  2B,  amorphous  boron  can  be  sep- 
arated with  some  difficulty  from  the  borides  of  magnesium  in  the 
resulting  mixture.  When  excess  of  powdered  aluminium  is  used,  hard 


638  INORGANIC  CHEMISTRY 

crystals  of  boron,  containing  aluminium,  are  found  in  the  solidified 
metal.  They  may  be  separated  by  interaction  of  the  metal  with 
dilute  hydrochloric  acid. 

Properties.  —  Amorphous  boron  is  a  black  powder.  It  unites 
with  the  same  elements  as  does  silicon  (p.  631),  but  with  somewhat 
greater  activity.  Like  carbon  (p.  425),  it  is  also  oxidized  by  hot,  con- 
centrated sulphuric  or  nitric  acid,  the  product  being  boric  acid.  The 
crystalline  variety  is  less  rapidly  attacked  in  each  case.  Both  kinds 
interact  with  fused  potassium  hydroxide,  giving  a  borate : 

2B  +  6KOH  -rf  2K3BO3  +  3H2. 

Hydrides  and  Halides  of  Boron.  —  When  magnesium  boride 
Mg3B2  is  treated  with  hydrochloric  acid,  a  gas  containing  much  hydro- 
gen, and  two  hydrides  of  boron,  B4Hi0  and  B2H6,  is  given  off. 

By  combined  action  of  carbon  and  chlorine  on  boric  oxide,  using 
the  principle  employed  in  preparing  silicon  tetrachloride  (p.  632),  the 
trichloride  of  boron  BC13  may  be  made.  It  is  likewise  formed  by 
direct  union  of  the  free  elements.  It  is  also  made  easily  by  heating 
boric  acid  and  phosphorus  pentachloride,  the  action  being  an  example 
of  the  behavior  of  the  latter  towards  hydroxyl  compounds  (cf.  p.  555) : 

B(OH)3  +  3PC15  ->  BC13  +  3POC13  +  3HC1. 

The  products  are  separated  by  fractional  distillation.  Boron  tri- 
chloride is  a  liquid  which  boils  at  18°,  fumes  strongly  in  moist  air, 
and  is  completely  hydrolyzed  by  water. 

Boron  trifluoride  BF3  is  made  by  the  interaction  of  calcium 
fluoride  and  sulphuric  acid  with  boron  trioxide.  The  mode  of 
preparation  and  the  properties  of  the  substance  recall  silicon  tetra- 
fluoride  (p.  632).  It  interacts  with  water,  like  the  latter,  giving 
boric  acid  and  hydrofluoboric  acid  HBF4: 

4BF3  +  3H20  -»  B(OH)3  +  3HBF4. 

The  boric  acid,  not  being  very  soluble,  is  precipitated.  Hydro- 
fluoboric acid  is  known  only  in  solution,  although  many  of  its  salts 
are  stable. 

Boric  Acid  H3BOS.  —  Boric  acid  (boracic  acid)  is  somewhat 
volatile  with  steam  (cf.  p.  563),  and  is  found  in  Tuscany  in  jets  of 
water  vapor  (soffioni)  which  issue  from  the  ground.  Water,  retained 


SILICON  AND  BORON  639 

in  small  basins  by  brickwork,  is  placed  over  the  openings,  and  from 
this  water,  after  evaporation  by  the  help  of  the  steam  of  the  soffioni 
themselves,  boric  acid  is  obtained  in  crystalline  form.  As  boric  acid 
is  a  very  feeble  acid,  and  withal  little  soluble,  it  may  be  made  by  mix- 
ing sulphuric  acid  and  concentrated  borax  solution,  and  cooling: 

NaJB407  +  H2S04  +  5H20  *=*  Na*S04  +  4H3B03  j. 

Boric  acid  crystallizes  from  water  in  thin  white  plates,  which  are 
soapy  (like  graphite  and  talc)  to  the  touch.  Its  solubility  in  water  is 
4  parts  in  100  at  19°  and  34  in  100  at  100°.  The  solution  scarcely 
affects  litmus.  The  green  tint  it  confers  on  the  Bunsen  flame  is  used 
as  a  test.  At  100°  the  acid  slowly  loses  water,  leaving  metaboric 
acid  HBO2,  and  at  140°  tetraboric  acid  is  formed:  4HBO2  -  H20 
— >  H2B4O7.  Strong  heating  gives  the  trioxide  B2O3.  When  dissolved 
in  water,  these  dehydrated  compounds  revert  to  boric  acid.  The 
solution  of  boric  acid  in  water  is  used  as  an  antiseptic  in  medicine 
(half -saturated,  2  per  cent  solution)  and  sometimes  as  a  preservative 
for  milk  and  other  foods. 

Borates.  —  Borates  derived  from  orthoboric  acid  are  practically 
unknown.  The  most  familiar  salt  is  borax  or  sodium  tetraborate. 
The  decahydrate  Na2B4O7,10H2O,  which  crystallizes  from  water  at 
27°  in  large,  transparent  prisms,  and  the  pentahydrate  which  crystal- 
lizes at  56°,*  are  both  marketed.  They  are  made  by  crystallization  of 
native  borax.  In  Germany,  borax  is  prepared  from  boracite,  found 
at  Stassfurt,  by  decomposing  a  solution  of  the  mineral  with  hydro- 
chloric acid: 

MgCl2,2Mg3B8Oi5  +  12HC1  +  18H2O  ->  7MgCl2  +  16B(OH)3. 

The  boric  acid  is  redissolved  in  boiling  water,  and  sodium  carbon- 
ate is  added:  4B(OH)3  +  Na^COs  ->  Na2B4O7  +  6H2O  +  CO2.  In 
California  it  is  made  from  colemanite  by  interaction  with  sodium 
carbonate: 

2CasB6On  +  4NasC03  +  H20  ->  4CaC03|  +  3Na2B407  +  2NaOH. 

Since  boric  acid  is  a  feeble  acid,  borax  is  hydrolyzed  by  water, 
and  the  solution  has  a  marked  alkaline  reaction  (c/.  p.  399).  In  a 
0.1  N  solution  (25°),  0.5  per  cent  is  hydrolyzed. 

*  For  explanation  of  the  relation  of  temperature  of  crystallization  to  degree  of 
hydration,  see  under  Manganous  sulphate. 


640  INORGANIC  CHEMISTRY 

When  heated  with  oxides  of  metals,  sodium  tetraborate  behaves 
like  sodium  metaphosphate  (cf.  p.  560),  and  is  used  in  the  form  of 
beads  in  analysis.  If  its  formula  be  written  2NaB02,B2O3,  it  will  be 
seen  that  a  considerable  excess  of  the  acid  anhydride  is  contained  in  it, 
and  that,  therefore,  a  mixed  metaborate  may  be  formed  by  union  with 
some  basic  oxide.  Thus,  with  a  trace  of  cupric  oxide,  the  bead  is 
tinged  with  green,  from  the  presence  of  a  compound  like  2NaBO2,- 
Cu(B02)2-  Cobalt  compounds  give  a  deep-blue  color  to  the  bead. 
For  the  same  reason,  borax  is  used  in  hard  soldering.  The  hard 
solder  (brass)  is  placed,  with  a  little  borax,  upon  the  joint  between  the 
objects  of  copper  or  brass  which  are  to  be  soldered.  At  the  tempera- 
ture produced  by  the  blast-lamp  the  borax  dissolves  the  superficial 
coating  of  oxides,  and  the  molten  solder  is  able  to  "wet"  the  clean 
surfaces.  In  welding  iron,  borax  is  scattered  on  the  parts,  and  com- 
bines with  the  oxide  to  form  a  fusible  mixed  borate,  which  is  forced 
out  by  the  pressure.  A  substance  used  thus,  to  bring  infusible  bodies 
into  a  fusible  form  of  combination,  is  called  a  flux.  Borax  is  also 
mixed  with  glass  in  making  enamels  for  cooking  utensils. 

Boron  Trioxide.  —  The  oxide,  as  made  by  heating  boric  acid,  is 
a  glassy  white  solid.  It  is  obtained  also  by  burning  boron  in  oxygen. 
Being  almost  perfectly  involatile,  it  is  able,  when  heated  with  salts,  to 
displace  other  acid  anhydrides  which  can  be  vaporized  : 

K2S04  +  B203  ^  2KB02  +  SO3T. 

It  has  a  slight  tendency  to  act  as  a  basic  oxide.  With  fuming  sul- 
phuric acid  it  gives  a  boryl  pyrosulphate  BO,HS2Or  which  is  decom- 
posed by  water,  and  with  phosphoric  acid  a  phosphate  BPO4  which  is 
stable. 

Nitride  and  Carbide.  —  One  of  the  difficulties  in  making  free 
boron  is  due  to  its  very  great  affinity  for  nitrogen,  with  which,  when 
heated  in  the  air,  it  unites  to  form  a  nitride  BN.  This  compound  is 
more  easily  made  by  heating  borax  with  ammonium  chloride  : 

4NH4C1  -»  4BN  +  2NaCl  +  7H2O  +  2HC1. 


The  nitride  is  a  white  solid  which  is  easily  decomposed  when  heated 
in  a  current  of  steam  (cf.  p.  514):  BN  +  3H2O->B(OH)3  +  NH3. 
A  carbide  of  boron  BeC  is  made  by  heating  the  free  substances  in 
the  electric  furnace.  It  is  harder  than  carborundum,  and  stands 
next  to  the  diamond  in  respect  to  hardness. 


SILICON  AND  BORON  641 

Exercises.  —  1.   Why  is  the  fact  that  carborundum  SiC  (p.  572) 
is  not  affected  by  water  or  acids  worthy  of  mention? 

2.  Compare  and  contrast  the  elements  carbon  and  silicon,  and 
their  corresponding  compounds. 

3.  Formulate  (p.  382)  the  interaction  between  aqueous  solutions 
of  an  ammonium  salt  and  of  sodium  orthosilicate. 


CHAPTER   XXXII 
THE  BASE-FORMING  ELEMENTS 

THERE  are  two  ways  in  which  the  chemistry  of  a  given  set  of  ele- 
ments may  be  described.  We  may  take  up  the  elements  in  succes- 
sion, and  discuss  under  each  its  physical  and  chemical  properties,  and 
the  manufacture  and  behavior  of  a  certain  number  of  its  compounds, 
such  as  the  oxide,  hydroxide,  nitrate,  and  sulphate.  Or  we  may 
arrange  our  major  classification  according  to  the  properties  and  the 
forms  of  combination,  and  detail  the  facts  about  the  same  set  of 
elements  under  each.  Both  methods  have  such  advantages  that 
neither  can  be  sacrificed  entirely.  We  .shall,  therefore,  adopt  the 
former  plan  for  our  division  into  chapters,  following  the  usage  already 
employed  for  the  non-metallic  elements.  In  the  present  chapter  a 
preliminary  view  of  the  chemistry  of  the  metals  will  be  given  according 
to  the  second  method. 

Physical  Properties  of  the  Metals.  —  A  knowledge  of  the 
physical  properties  of  the  metals  is,  to  the  chemist,  of  the  greatest 
importance  in  connection  with  their  manufacture  and  treatment. 
The  following  brief  statement  in  regard  to  some  of  these  properties  is 
illustrative  rather  than  exhaustive.  It  should  be  noticed  that  the 
properties  of  a  metal  vary  according  as  the  specimen  has  been  pre- 
pared by  rolling,  casting,  or  some  other  process.  Numerical  values, 
therefore,  when  given,  are  only  approximate. 

Metals  show  what  is  commonly  called  a  metallic  luster,  but,  as  a 
rule,  they  do  so  only  when  in  compact  form.  Magnesium  and  alumin- 
ium exhibit  it  when  powdered,  but,  in  this  condition  most  of  the 
metals  are  black.  In  compact  masses  the  metals  are  usually  silvery 
white  in  color.  Gold  and  copper,  which  are  yellow  and  red  respec- 
tively, are  the  conspicuous  exceptions. 

The  metals  can  all  be  obtained  in  crystallized  form,  when  a  fused 
mass  is  allowed  to  cool  slowly  and  the  unsolidified  portion  is  poured 
off.  In  almost  all  cases  the  crystals  belong  to  the  regular  system. 
With  the  metals  most  nearly  allied  to  the  non-metals,  however,  they 
do  not.  Thus,  the  crystals  of  antimony  and  bismuth  belong  to  the 
hexagonal  system,  and  those  of  tin  to  the  square  prismatic. 

642 


THE  BASE-FORMING  ELEMENTS 


643 


The  metals  vary  in  density  from  lithium,  which  is  little  more  than 
half  as  heavy  as  water  (sp.  gr.  0.53),  to  osmium,  the  density  of  which 
is  22.5.  Those  which  have  a  density  less  than  5,  namely,  potassium, 
sodium,  calcium,  magnesium,  aluminium,  and  barium,  are  called  the 
light  metals,  and  the  others  the  heavy  metals. 

Most  metals  are  malleable,  and  can  be  beaten  into  thin  sheets  with- 
out loss  of  continuity.  Those  which  are  allied  to  the  non-metals, 
however,  such  as  arsenic,  antimony,  and  bismuth,  are  brittle,  and  can 
be  reduced  to  powder  in  a  mortar.  Zinc  becomes  malleable  only 
when  heated  to  150°.  The  order  of  the  elements  in  respect  to  this 
property,  beginning  with  the  most  malleable,  is :  Au,  Ag,  Cu,  Sn,  Pt, 
Pb,  Zn,  Fe,  Ni. 

The  tenacity  of  the  metals  places  them  in  a  different  order. 
It  is  measured  by  the  number  of  kilograms  which  a  piece  of  the  metal 
1  sq.  mm.  in  section  can  sustain  without  breaking.  The  values  are 
as  follows:  Fe  62,  Cu  42,  Pt  34,  Ag  29,  Au  27,  Al  20,  Zn  5,  Pb  2. 

The  hardness  (Appendix  II)  is  measured  by  the  ease  with  which  the 
material  may  be  disintegrated  by  a  sharp,  hard  instrument.  Potas- 
sium is  as  soft  as  wax,  while  chromium  is  hard  enough  to  cut  glass. 

The  temperature  at  which  the  metal  fuses  has  an  important 
bearing  on  its  manufacture.  Most  of  the  following  melting-points 
are  only  approximate: 


Mercury  .... 
Potassium  .  .  . 
Sodium  .... 

-39° 
62° 
97° 

Zinc  
Antimony  .  .  . 
Magnesium  .  . 

419° 
630° 
651° 

Cast  iron  .  .  . 
Manganese  .  . 
Nickel  ..... 

1150° 
1260° 
1452° 

Tin  
Bismuth  .  .  . 

232° 
271° 

Aluminium  .  . 
Silver 

659° 
960° 

Chromium  .  .  . 
Iron  (pure) 

1520° 
1530° 

Cadmium  .  .  . 

321° 

Gold 

1063° 

Platinum 

1755° 

Lead  

327° 

Copper  .... 

1083° 

Tungsten  .  .  . 

3540° 

It  will  be  seen  that  mercury  is  a  liquid,  that  potassium  and  sodium 
melt  below  the  boiling-point  of  water,  and  that  the  metals  down  to  the 
foot  of  the  second  column  can  be  melted  easily  with  the  Bunsen 
flame.  The  metals  osmium,  molybdenum,  uranium,  iridium,  and 
vanadium  have  melting-points  above  that  of  platinum.  For  iron, 
two  different  methods  give  1510°  and  1550°,  respectively. 

The  methods  of  manufacture  and  the  treatment  of  metals  are 
much  influenced  also  by  their  volatility.  The  following  are  easily 
distilled:  Mercury,  b.-p.  357°;  potassium  and  sodium,  b.-p.  about 
700°;  cadmium,  b.-p.  770°;  zinc,  b.-p.  920°.  Even  the  most  in  vola- 
tile metals  can  be  converted  into  vapor  in  the  electric  arc. 


644  INORGANIC  CHEMISTRY 

In  many  cases  molten  metals  dissolve  in  one  another  freely.  The 
mixtures  are  called  alloys,  and  in  some  cases  are  simply  solid  solu- 
tions. Sometimes,  as  in  the  case  of  lead  and  tin,  mixtures  can  be 
formed  in  all  proportions.  Oh  the  other  hand,  the  solubility  may  be 
limited,  as  in  the  case  of  zinc  and  lead,  where  only  1.6  parts  of  the 
former  dissolve  in  100  parts  of  the  latter.  Frequently  chemical 
compounds  are  formed.  The  colors  of  alloys  are  not  the  average  of 
those  of  the  constituents.  Thus,  a  mixture  containing  copper  with 
30  per  cent  of  tin  is  perfectly  white.  A  similar  mixture  with  30 
per  cent  of  zinc  is  pale  yellow.  The  nickel  alloy  used  in  coining 
contains  75  per  cent  of  copper  and  25  per  cent  of  nickel,  yet  it  shows 
none  of  the  color  of  the  former.  Thirty  per  cent  of  gold  may  be 
added  to  silver  without  conferring  any  yellow  tint  upon  it. 

Some  of  the  properties  of  alloys  are  classifiable  by  the  ordinary 
laws  of  solution.  Thus,  a  foreign  solute  lowers  the  vapor  tension  of  a 
solvent  (p.  197),  and  so  the  presence  of  a  foreign  metal  diminishes  the 
ease  with  which  particles  can  be  torn  from  the  surface  by  any  means 
whatever.  That  is  to  say,  it  increases  the  hardness.  A  foreign  metal 
also  lowers  the  melting-point  (p.  199,  see  solder).  In  many  cases  the 
metal  becomes  less  active  when  alloyed.  Thus,  a  mixture  of  gold  and 
silver  containing  twenty-five  per  cent  of  the  latter  does  not  interact 
visibly  with  nitric  acid.  It  is  necessary  to  bring  the  amount  of  silver 
up  to  75  per  cent  at  least  ("quartation")  in  order  that  the  silver  may 
be  freely  attacked  by  the  warm  acid.  The  gold  remains  in  any  case 
untouched.  The  malleability  and  the  conductivity  for  heat  and  elec- 
tricity are  diminished  by  solution  of  a  foreign  metal.  Copper,  whose 
commercial  applications  depend  largely  on  the  first  and  third  of  these 
properties,  is  much  affected  in  respect  to  them  by  the  presence  of  even 
small  traces  of  impurities. 

Alloys  in  which  mercury  forms  one  of  the  components  are  known 
as  amalgams  (Gk.  //.aXay/ia,  a  soft  mass) ,  and  are  formed  with  especial 
ease  by  the  lighter  metals.  Of  the  common  metals,  iron  is  the  least 
miscible  with  mercury. 

The  good  conductivity  of  metals  for  electricity  distinguishes  them 
with  some  degree  of  sharpness  from  the  non-metals.  They  show  con- 
siderable variation  amongst  themselves,  silver  conducting  sixty  times 
as  well  as  mercury.  The  conductivity  increases  as  the  temperature  is 
lowered,  and  this  fact  is  taken  advantage  of  in  the  measurement  of 
the  temperature  of  liquefied  gases.  The  platinum  resistance  ther- 
mometer consists  chiefly  of  a  wire  of  platinum.  The  resistance  of 
this  metal  diminishes  so  rapidly,  with  decreasing  temperature,  that 


THE  BASE-FORMING  ELEMENTS  645 

measurement  of  its  resistance  can  be  used  for  the  accurate  determina- 
tion of  the  temperature.  In  the  following  table  the  conductivities 
of  the  metals  are  expressed  in  terms  of  the  number  of  meters  of  wire 
1  sq.  mm.  in  section  which,  at  15°,  offer  a  resistance  of  one  ohm: 

Silver,  cast 62.89  Nickel,  cast  .    .    .    .  7.59 

Copper,  commercial      .    .    .57.40  Iron,  drawn  .    .    .    .7.55 

Gold,  cast 46.30  Platinum 5.7-8.4 

Aluminium,  commercial  .    .   31.52  Steel 5.43 

Zinc,  rolled 16.95  Lead 4.56 

Brass 14.17  Mercury 1.049 

The  resistance  at  0°  of  a  column  of  mercury  1  sq.  mm.  in  section 
and  1.063  meters  long  is  the  unit  of  resistance,  and  is  called  one  ohm. 
It  is  employed  in  expressing  the  conductivities  of  solutions  (p.  362). 
To  compare  the  above  figures  with  those  given  for  solutions,  however, 
it  must  be  recalled  that,  in  the  measurement  of  the  conductivities 
of  the  latter,  a  column  only  1  cm.  in  length  and  of  1  sq.  cm.  area  was 
employed,  so  that  the  figures  representing  the  conductivities  of 
solutions  are  on  a  scale  approximately  ten  thousand  times  as  great 
as  those  presented  in  the  above  table.  Thus,  normal  hydrochloric 
acid  (p.  365)  has  a  conductivity  on  the  above  scale  of  0.0301,  or  less 
than  a  thirtieth  of  that  of  mercury. 

The  world's  production  (1913)  of  the  metals  in  metric  tons  of  1000 
kilos,  is  approximately  as  follows: 


Copper          1,000,000 

Chromium 

50,000 

Gold 

680 

Zinc               1,000,000 

Nickel 

32,000 

Bismuth 

500 

Lead              1,000,000 

Silver 

7,800 

Cadmium 

50 

Tin                   120,000 

Tungsten 

4,800 

Platinum 

9 

Aluminium        79,000 

Mercury 

3,000 

General  Chemical  Relations  of  the  Metallic  Elements.  — 

Since  most  of  the  compounds  of  the  metals  are  ionogens,  their  solu- 
tions, except  when  the  metal  is  a  part  of  a  compound  ion  or  of  a  com- 
plex ion  (see  below),  all  contain  the  metal  in  the  ionic  state,  and  the 
resulting  substances,  such  as  potassium-ion  and  cupric-ion,  have 
constant  properties,  irrespective  of  the  nature  of  the  negative  ion 
with  which  they  may  be  mixed.  The  properties  of  the  ions,  simple 
and  compound,  are  much  used  in  making  tests  in  analytical  chem- 
istry. On  the  other  hand,  the  chemical  properties  of  the  oxides  and 
of  the  salts  in  the  dry  state  are  of  importance  in  connection  with 
metallurgy. 


646  INORGANIC  CHEMISTRY 

There  are  three  chemical  relations  which  are  characteristic  of  all 
metallic  elements.  These  form  the  basis  of  the  distinction  between 
metallic  and  non-metallic  elements.  The  first  two  of  them  have 
already  been  discussed  somewhat  fully. 

1.  The  metals  are  able  by  themselves  to  form  positive  radicals  of 
salts,  and,  therefore,  to  exist  alone  as  positive  ions  (pp.  356,  375). 

2.  The  oxides  and  hydroxides  of  the  metal  sare  basic  (pp.  149, 374). 

3.  Each  typical  metal  has  at  least  one  halogen  compound  which 
is  little,  if  at  all,  hydrolyzed  by  water  (p.  399  and  next  section). 
The  same  thing  is  true  of  nitrates  and  other  salts  involving  active 
acids. 

There  are  three  other  chemical  relations  which  are  shown  by 
many  metallic  elements,  but  no  one  of  them  applies  to  all. 

4.  An  oxide  or  hydroxide  which  is  basic  may  also  be  acidic,  as, 
for  example,  zinc  hydroxide  (p.  122).     Even  when  this  is  not  the 
case,  some  other  oxide  of  the  metal  may  be  acidic  exclusively,  as  is 
manganese  heptoxide  (p.  457).     In  consequence  of  either  of  these 
facts,  a  metal  may  form  part  of  the  negative  radical  of  a  simple  salt, 
and  therefore  be  found  in  a  negative  ion,  as,  HZnO2~  or  MnO4~. 

5.  Some  salts  of  certain  metals  combine  with  those  of  others  to 
give  complex  salts  (p.  405  and  see  p.  649).     Of  this  sort  are  the 
complex  cyanides,  such  as  K.Ag(CN)2  and  K4.Fe(CN)e.     A  metal 
thus  forms  part  of  the  negative  radical  of  a  salt  of  a  complex  acid, 
and  therefore  is  found  in  an  anion  like  Ag(CN)2~  or  PtCle". 

6.  Some  metals  also  form  parts  of  complex  cations  which  are  con- 
tained in  solutions  of  molecular  compounds  (p.  530).     Thus,  when 
AgCl,3NH3  is  dissolved  in  water,  or  when  ammonium  hydroxide  is 
added  to  a  solution  of  a  salt  of  silver,  the  positive  ion  is  found  to  bs 
Ag(NH3)2+  (see  under  Copper).     (For  a  detailed  illustration  of  the 
application  of  these  six  criteria,  see  discussion  of  the  chemical  rela- 
tions in  the  nitrogen  family,  Chap.  XLI.) 

Aside  from  these  points,  many  features  in  the  behavior  of  metals 
and  their  compounds  are  summed  up  in  the  electromotive  series  (p. 
404).  The  reader  should  re-read  all  the  parts  referred  to  above 
before  proceeding  farther.  He  should  also  reexamine  the  various 
kinds  of  chemical  changes  enumerated  on  p.  228  and  particularly  the 
varieties  of  ionic  chemical  change  on  p.  402. 

Hydrolysis  of  Halogen  Compounds,  Used  to  Distinguish 
Metallic  Elements  from  Non-Metallic  Elements.  —  We  have 
seen  that  the  halogen  compounds  of  phosphorus  (p.  555),  of  sulphur 


THE   BASE-FORMING  ELEMENTS  647 

(p.  450),  of  silicon  (p.  632),  and  of  other  non-metallic  elements,  are 
completely  hydrolyzed  by  water,  giving  the  hydrogen  halide,  and  an 
acid  which  contains  the  hydroxyl  of  the  water: 

PC13  +  3HOH  ->  3HC1  +  P(OH)3. 

Now,  those  elements  whose  halogen  compounds  are  not  hydrolyzed  by 
water,  or,  at  all  events,  are  only  partly  hydrolyzed,  are  the  ones 
classed  as  metallic  elements.  Thus,  sodium  chloride  is  not  decom- 
posed appreciably  by  water,  and  cupric  chloride,  like  cupric  sulphate 
(p.  399),  is  but  slightly  hydrolyzed,  and  its  solution  has  a  faint  acid 
reaction : 

CuCl2  +  2H20  <=*  2HC1  +  Cu(OH)2, 
or,  Cu++  +  2OH~  +  2H+  ->  Cu(OH)2  +  2H+. 

In  a  few  cases,  as  with  the  chlorides  of  antimony  (q.v.)  and  of  bis- 
muth (q.v.),  a  considerable  proportion,  but  not  all,  of  the  halogen  is 
removed  from  each  molecule: 

SbCl3  +  H20  <=±  2HC1  +  SbOClJ . 

Here  the  water  produces  a  slight  hydrolysis,  but  the  insolubility  of 
one  product  weakens  the  reverse  action  and  the  equilibrium  is  dis- 
placed forwards  chiefly  because  of  the  precipitation.  The  resulting 
mixture  is  strongly  acid,  and  the  product  (antimony  oxychloride)  is 
a  definite  compound,  of  the  nature  of  a  mixed  salt  (p.  401),  known  as 
a  basic  salt.  The  difference  is  that,  with  the  halides  of  the  metallic 
elements,  the  action  on  water  is  notably  reversible,  and  the  reverse 
action,  unless  handicapped  by  precipitation  of  one  factor,  prevents 
much  displacement  of  the  equilibrium,  while  with  halides  of  the 
non-metallic  elements,  the  action  is  not  reversible. 

Hydrolysis  of  the  halides  of  the  metals  is  increased  by  rise  in  tem- 
perature and  by  dilution  of  the  solution  (addition  of  more  water), 
and  also  gains  headway  when  one  of  the  products  of  hydrolysis  is 
thrown  down  as  a  precipitate.  The  last  two  influences  are  the  ones 
which  normally  permit  any  reversible  action  to  approach  completion 
(p.  301). 

The  halogen  compounds  are  chosen  as  the  basis  of  this  criterion 
because  the  halogen  acids  are  active  and  would  reverse  the  hydrolysis 
completely,  and  leave  no  acid  reaction,  if  the  result  depended  upon 
them  alone.  It  is  the  lack  of  activity  in  the  base,  and  the  tendency 
of  its  molecules  to  be  formed  from  the  metal  ions  of  the  salt  and  the 
OH~  of  the  water  (p.  399),  .that  determine  the  slight  hydrolysis, 


648  INORGANIC  CHEMISTRY 

when  it  occurs.  Thus,  this  criterion  is  simply  another  means  of  recog- 
nizing whether  or  not  the  hydroxide  of  the  element  is  a  strong  (much 
ionized)  base,  and  its  application  gives,  therefore,  the  same  result  in 
each  case  as  does  the  employment  of  the  second  of  the  chemical  char- 
acteristics of  the  metallic  elements  (see  above). 

Other,  non-halide  salts  of  the  metals,  even  of  the  most  active, 
may  be  extensively  hydrolyzed  by  water.  Thus,  sodium  sulphide 
is  decomposed  by  it  (p.  418)  to  the  extent  of  one-half.  But  here  the 
solution  is  alkaline  in  reaction: 

Na2S  +  H20  ->  NaSH  +  NaOH, 

owing  to  the  small  ionization  of  the  SH~  ion,  and  the  result  is  due 
to  the  feebleness  of  the  acid  H2S.  Indeed,  the  great  activity 
of  the  base  is  demonstrated  by  the  final  reaction  of  the  solu- 
tion, and  the  metallic  nature  of  sodium  is  therefore  not  impugned 
by  the  existence  of  hydrolysis  per  se  in  such  a  salt  as  this,  but  is 
rather  confirmed. 

To  sum  up:  An  alkaline  reaction  shows  that  we  have  a  solution  of 
a  salt  of  an  active  metal  with  a  weak  acid;  an  acid  reaction  that  we 
have  a  salt  of  an  inactive  metal,  which  may  even  verge  on  the  non- 
metallic,  with  an  active  acid.  Salts  of  two  active  components  give 
neutral  solutions.  Thus,  as  a  particular  case,  the  halogen  com- 
pounds of  the  typical  metals  are  not  perceptibly  hydrolyzed  by  water, 
and  the  hydrolysis  of  the  halide  of  a  less  pronounced  metal  takes 
place  with  acid  reaction,  and  is  easily  reversible  by  excess  of  either 
product.  A  salt  of  an  acid  and  base  both  of  which  are  weak  is  also 
hydrolyzed.  If  the  resulting  base  or  acid  is  insoluble,  the  hydrolysis 
may  go  nearly  to  completion.  Aluminium  carbonate  and  ammonium 
silicate  are  examples  of  salts  which,  for  this  reason,  are  completely 
hydrolyzed.  From  the  former,  aluminium  hydroxide  is  precipitated, 
and  from  the  latter  silicic  acid.  When  both  base  and  acid  are  weak, 
but  soluble,  the  resulting  mixture  may  have  an  acid  or  a  basic  reaction, 
if  the  acid  or  the  base  is  sufficiently  active  to  affect  an  indicator. 
Thus,  ammonium  sulphide  (NH4)2S  solution  is  alkaline. 

The  rather  exaggerated  language  commonly  used  by  chemists  in 
regard  to  hydrolysis  must  not  be  misinterpreted.  When  it  is  said  that 
borax  gives  a  strongly  alkaline  solution,  and  is  "extensively"  hydro- 
lyzed, this  only  means  that  in  a  0.1  AT"  solution  0.5  per  cent  of  the  salt 
is  decomposed.  Aluminium  chloride  gives  a  strongly  acid  solution, 
but  even  in  a  0.001N  solution  the  hydrolysis  (25°)  reaches  only  4.5 
per  cent. 


THE  BASE-FORMING  ELEMENTS  649 

Salts  of  Complex  Acids.  —  These  salts  are  of  many  classes, 
and  arise  by  direct  union  of  two  salts.  Thus  we  have  cyanides  like 
potassium  argenticyanide  K.Ag(CN)2,  potassium  cuprocyanide 
K.Cu(CN)2,  and  potassium  ferrocyanide  K4.Fe(CN)e: 

K.CN  +  AgCNT  ->  K.Ag(CN)2. 

The  complex  sulphur  compounds  of  arsenic,  antimony,  and  tin  (q.v.), 
such  as  sodium  sulphostannate  Na2.SnS3,  are  made  in  the  same  way  : 

SnS2T«=±Na2.SnS3. 


Many  double  halides  are  of  a  like  nature,  as  potassium  chloroplat- 
inate  K^.PtCle,  chloroplatinic  acid  H^.PtCle,  and  sodium  chloraurate 
Na.AuCLi.  In  every  case,  the  metal  of  one  of  the  original  salts  is 
contained  in  the  negative  radical  (the  anion).  Hydrofluosilicic  acid 
(p.  633)  is  a  compound  analogous  to  those  of  the  last  group,  excepting 
that  SiF4,  from  which  it  is  formed,  is  not  a  salt,  and  silicon  is  not  a 
metal. 

The  characteristic  of  these  compounds  is  that,  when  they  are 
ionized,  the  less  positive  metal  is  contained  in  a  complex  anion  like 
Ag(CN)2~,  SnS3=,  PtCl6=.  In  fact,  they  behave,  in  most  respects, 
like  ordinary  single  salts.  Thus,  they  undergo  double  decomposition 
in  the  normal  manner,  the  complex  ion  acting  as  a  whole.  For 
example,  a  soluble  ferrocyanide  with  a  zinc  salt  gives  zinc  ferro- 
cyanide : 

K4.Fe(CN)6  +  2Zn.S04  +±  Zn2Fe(CN)6|  +  2K2.S04. 

This  behavior  distinguishes  salts  of  complex  acids  from  double  salts 
(p.  402).  In  typical  cases  the  latter  resemble  the  former,  indeed, 
only  in  then-  mode  of  preparation  (by  union  of  two  simple  salts),  and 
in  solution  are  resolved  once  more  into  the  component  salts  and  their 
separate  ions.  The  distinction  must  be  regarded,  however,  only  as  a 
means  of  rough  classification  for  practical  purposes.  In  reality,  all 
complex  ions  give  at  least  a  trace  of  the  simpler  ions,  and  many  are  de- 
composed to  a  noticeable  extent.  At  the  other  extreme,  the  double 
salts  form  complex  ions  in  appreciable  amounts  in  concentrated 
solutions. 

In  harmony  with  the  above  characteristic,  salts  of  complex  acids, 
like  potassium  ferrocyanide,  when  treated  with  acids,  like  sulphuric 
acid,  undergo  double  decomposition,  and  give  the  free  complex  acid 
(here  hydrof  errocyanic  acid)  : 

K4.Fe(CN)6  +  2H2.S04  +±  H4Fe(CN)6  1  +  2K2.S04. 


650  INORGANIC   CHEMISTRY 

But  in  many  cases  the  acids  are  unstable,  those  of  the  cyanides,  for 
example,  giving  up  hydrocyanic  acid  and  those  of  the  complex  sulphur 
compounds  giving  hydrogen  sulphide  : 

+SnS2|. 


The  complex  halogen  acids,  however,  like  H^.PtCle  and  H.AuCU,  are 
stable,  and  are  fairly  active  as  acids. 

The  similarity  of  these  compounds  to  oxygen  acids,  and  their  salts,  may  be 
seen  if  we  imagine  them  to  be  cases  in  which  cyanogen,  sulphur,  chlorine,  and  other 
radicals  have  taken  the  place  of  oxygen  in  the  anion.  Thus,  [(CN)2]11,  [Cl2]n,  and 
S11  are  equivalent  to  O11.  Many  of  the  corresponding  oxygen  compounds  are 
actually  known,  as  sodium  stannate  Na^.SnOs  and  sodium  aurate  Na.AuO2,  corre- 
sponding respectively  to  Naa.SnSa  and  Na.AuCU. 

The  behavior  of  complex  ions  is  discussed  further  under  Copper. 

Classification  of  the  Metallic  Elements  by  their  Chemical 
Relations.  —  In  treating  of  the  metallic  elements  and  their  com- 
pounds we  shall  use  the  groupings  provided  by  the  periodic  system 
(p.  461).  A  division  into  eleven  sets,  which  are  described  briefly  and 
in  general  terms  below,  will  be  sufficient  for  the  purpose  of  this 
book: 

1.  Metals  of  the  Alkalies.  —  Lithium,  sodium,  potassium,  rubid- 
ium, caesium,  and  the  radical  ammonium  NH4.     These  metals  are 
univalent,  and  their  oxides  and  hydroxides  have  strongly  basic  prop- 
erties.    Their  salts  with  active  acids  are  not  hydrolyzed  in  solution. 

2.  Metals  of  the  Alkaline  Earths.  —  Calcium,  strontium,  barium, 
and  radium.     These  metals  are  bivalent  in  all  their  compounds. 
Their  oxides  and  hydroxides  are  strongly  basic,  but  the  latter  are 
not  so  soluble  in  water  as  are  the  hydroxides  of  the  former  family. 
The  salts  with  active  acids  are  not  hydrolyzed. 

3.  Copper,    Silver,    and   Gold.  —  These  metals  occupy  the  right 
side  of  the  second  column  of  the  table  (opposite  rear  cover)  .     Their 
alliance  with  the  alkali  metals,  their  neighbors,  is  remote.     Each, 
however,  gives  compounds  in  which  it  is  univalent,  although,  in 
their  commoner  compounds,  copper  and  gold  are  bivalent  and  triva- 
lent,  respectively.     The  oxides  and  hydroxides  of  copper  and  gold 
have  rather  weak  basic  properties;    those  of  silver  are  much  more 
active. 

4.  Beryllium,  Magnesium,  Zinc,  Cadmium,  and  Mercury.  —  These 
metals,  occupying  the  right-hand  side  of  the  third  column,  are  biva- 
lent, although  mercury  forms  a  series  of  compounds  in  which  it  is 


THE  BASE-FORMING  ELEMENTS  651 

univalent  as  well.    The  oxides  and  hydroxides  are  feebly  basic,  those 
of  magnesium  being  the  most  basic. 

5.  Aluminium  and  the  other  metals  on  both  sides  of  the  fourth 
column.  —  The  metals  of  these  groups  are  trivalent,  and  the  oxides 
and  hydroxides  are  feebly  basic  in  character.     The  hydroxide  of 
aluminium  has  also  a  feebly  acidic  tendency. 

6.  Germanium,  Tin,  and  Lead,  and  the  Titanium  Family. —  In  ac- 
cordance with  their  position  in  the  periodic  table  these  metals  are  all 
quadrivalent.     At  the  same  time  they  act  also  as  bivalent  elements, 
the  compounds  of  this  class  in  the  case  of  lead  being  the  more  familiar. 
The  oxides  and  hydroxides  are  feebly  basic,  and  are  able  also  to  play 
the  role  of  acidic  oxides  towards  strong  bases. 

7.  Arsenic,  Antimony,  and  Bismuth,  and  the  metals  of  the  Vana- 
dium Group.  —  These  elements,  like  nitrogen  and  phosphorus,  form 
two  sets  of  compounds  in  which  they  are  trivalent  and  quinquivalent, 
respectively.     The  acidic  tendency  of  the  oxides  and  hydroxides, 
which  in  the  last  column  was  noticeable,  is  here  much  more  pro- 
nounced.    Both  oxides  of  arsenic  are  almost  wholly  acidic  in  be- 
havior, and  the  pentoxides  of  the  other  elements  are  likewise  acid 
anhydrides  exclusively.     The  trioxides  are  basic  in  a  feeble  way, 
and  their  salts  are  much  hydrolyzed  by  water. 

8.  Chromium,    Molybdenum,  Tungsten,   and  Uranium.  —  These 
elements,  occupying  the  left-hand  side  of  the  seventh  column,  exhibit 
a  considerable  variety  of  valence.     The  maximum,  however,  is  six  in 
each  case.     The  oxides  of  the  form  CrO  and  Cr2O3,  in  which  the 
elements  are  bivalent  and  trivalent,  are  base-forming.     Those  of  the 
form  CrOs,  in  which  the  elements  are  sexivalent,  resemble  sulphur 
trioxide  and  are  acid  anhydrides. 

9.  Manganese.  —  This  element,  the  only  one  in  the  eighth  column 
which  has  not  yet  been  treated,  gives  several  series  of  compounds 
in  which  its  valance  varies  from  2  to  7.     Compounds  derived  from 
the  basic  oxide  MnO  are  the  salts  in  which  manganese  is  most  dis- 
tinctly a  metallic  element.    The  highest  oxide,  Mn2O7,  is  an  acid 
anhydride. 

10.  Iron,  Nickel,  and  Cobalt.  —  These  elements  give  oxides  which 
are  feebly  basic.     Iron  gives  two  extensive  series  of  compounds  in 
which  it  is  bivalent  and  trivalent,  respectively.     Those  of  the  former 
set  resemble  the  bivalent  salts  of  manganese  and  zinc.     Those  of  the 
latter  resemble  the  salts  of  aluminium.     Cobalt  and  nickel  in  most  of 
their  compounds  are  bivalent  elements,  and  the  behavior  recalls  that 
of  the  compounds  of  bivalent  manganese  and  zinc. 


652  INORGANIC  CHEMISTRY 

11.  Palladium  and  Platinum  Families.  —  The  metals  of  these  fam- 
ilies have  little  chemical  activity,  and  their  compounds  are  easily 
decomposed  by  heating.  Along  with  gold,  silver,  and  mercury,  which 
have  similar  characteristics,  they  are  sometimes  grouped  together 
under  the  name  of  the  noble  metals. 

Occurrence  of  the  Metals  in  Nature.  —  The  minerals  from 
which  metals  are  extracted  are  known  as  ores.  They  present  a  com- 
paratively small  number  of  different  kinds  of  compounds.  Most  of 
the  metals  are  found  in  more  than  one  of  these  forms,  so  that  in  the 
following  statement  the  same  metal  frequently  occurs  more  than 
once. 

When  the  metal  occurs  free  in  nature  it  is  said  to  be  native.  Thus 
we  have  native  gold,  silver,  metals  of  the  platinum  group,  copper, 
mercury,  bismuth,  antimony,  and  arsenic  (cf.  p.  404). 

The  metals  whose  oxides  are  important  minerals  are  iron,  man- 
ganese, tin,  zinc,  copper,  and  aluminium.  The  metals  are  obtained 
commercially  from  the  oxides  in  each  of  these  cases. 

The  metals  whose  sulphides  are  used  as  ores  are  iron,  nickel, 
cobalt,  antimony,  lead,  cadmium,  zinc,  and  copper. 

From  the  carbonates  we  obtain  iron,  lead,  zinc,  and  copper. 
Several  other  metals,  such  as  manganese,  magnesium,  barium, 
strontium,  and  calcium  occur  in  larger  or  smaller  quantities  in  the 
same  form  of  combination. 

The  metals  which  occur  as  sulphates  are  those  whose  sulphates 
are  not  freely  soluble,  namely,  lead,  barium,  strontium,  and  calcium. 

Compounds  of  metals  with  the  halogens  are  not  so  numerous. 
Silver  chloride  furnishes  a  limited  amount  of  silver.  Sodium  and  po- 
tassium chlorides  are  found  in  the  salt-beds,  and  cryolite  3NaF,AlF3 
is  used  in  the  manufacture  of  aluminium. 

The  natural  silicates  are  very  numerous,  but  are  seldom  used  for 
the  preparation  of  the  metals  (see,  however,  zinc).  Many  of  them 
are  employed  for  other  commercial  purposes,  kaolin  (p.  636)  or  clay 
being  a  conspicuous  example  of  this  class. 

Methods  of  Extraction  from  the  Ores.  —  The  art  of  extract- 
ing metals  from  their  ores  is  called  metallurgy.  Where  the  metal  is 
native,  the  process  is  simple,  since  melting  away  from  the  matrix 
(p.  410)  is  all  that  is  required.  Frequently  a  flux  (p.  640)  is  added. 
A  flux  usually  is  a  substance  which  interacts  with  infusible  materials 
to  give  fusible  ones.  It  combines  with  the  matrix,  giving  a  fusible 


THE  BASE-FORMING  ELEMENTS  653 

slag  (resembling  glass).  Since  the  slag  is  a  melted  salt,  usually  a 
silicate,  and  does  not  mix  at  all  with  the  molten  metal,  separation  of 
of  the  products  is  easily  effected.  When  the  ore  is  a  compound, 
the  metal  has  to  be  liberated  by  our  furnishing  a  material  capable  of 
combining  with  the  other  constituent.  The  details  of  the  process 
depend  on  various  circumstances.  Thus  the  volatile  metals,  like 
zinc  and  mercury,  are  driven  off  in  the  form  of  vapor,  and  secured 
by  condensation.  The  involatile  metals,  like  copper  and  iron,  run 
to  the  bottom  of  the  furnace  and  are  tapped  off. 

Where  the  ore  is  an  oxide  it  is  usually  reduced  by  heating  with 
carbon  in  some  form.  This  holds  for  the  oxides  of  iron  and  copper, 
for  example.  Some  oxides  are  not  reducible  by  carbon  in  an  ordinary 
furnace.  Such  are  the  oxides  of  calcium,  strontium,  barium,  mag- 
nesium, aluminium,  and  the  members  of  the  chromium  group.  At 
the  temperature  of  the  electric  furnace  even  these  may  be  reduced, 
but  the  carbides  are  formed  under  such  circumstances,  and  the  metals 
are  more  easily  obtained  otherwise.  Recently,  heating  the  pul- 
verized oxide  with  finely  powdered  aluminium  has  come  into  use, 
particularly  for  operations  on  a  small  scale.  Iron  oxide  is  easily 
reduced  by  this  means,  and  even  the  metals  manganese  and  chro- 
mium may  be  liberated  from  their  oxides  quite  readily  by  this  action. 
This  procedure  has  received  the  name  aluminotherxny  (q.v.),  on  ac- 
count of  the  great  amounts  of  heat  liberated.  In  the  laboratory 
the  oxides  of  the  less  active  metals  are  frequently  reduced  in  a  stream 
of  hydrogen  (cf.  p.  127). 

When  the  ore  is  a  carbonate,  it  is  first  heated  strongly  to  drive  out 
the  carbon  dioxide  (cf.  p.  573) :  FeCO3  ?±  FeO  +  CO2  f,  and  then  the 
oxide  is  treated  according  to  one  of  the  above  methods.  When  the 
ore  is  a  sulphide,  it  has  to  be  calcined  (cf.  p.  424)  in  order  to  remove 
the  sulphur,  and  the  resulting  oxide  is  then  treated  as  described  above. 

Chlorides  and  fluorides  of  the  metals  can  be  decomposed  by  heat- 
ing with  metallic  sodium  (cf.  p.  630).  This  method  was  formerly  em- 
ployed in  the  making  of  magnesium  and  aluminium. 

The  metals  which  are  not  readily  secured  in  any  of  the  above  ways, 
can  be  obtained  easily  by  electrolysis  of  the  fused  chloride  or  of 
some  other  simple  compound.  Aluminium  is  now  manufactured 
entirely  by  the  electrolysis  of  a  solution  of  aluminium  oxide  in  mol- 
ten cryolite. 

Compounds  of  the  Metals:   Oxides  and  Hydroxides.  —  The 

oxides  may  be  made  by  direct  burning  of  the  metals,  by  heating  the 


654  INORGANIC  CHEMISTRY 

nitrates  (cf.  p.  531),  the  carbonates  (cf.  p.  573),  or  the  hydroxides: 
Ca(OH)2  ^  CaO  +  H2Ot-  They  are  practically  insoluble  in  water, 
although  those  of  the  metals  of  the  alkalies  and  of  the  metals  of  the 
alkaline  earths  interact  with  water  rapidly  to  give  the  hydroxides. 
They  are  usually  stable.  Those  of  gold,  platinum,  silver,  and  mer- 
cury, decompose  when  heated,  yet  with  increasing  difficulty  in  this 
order.  The  metals,  like  the  non-metals,  frequently  give  several 
different  oxides.  Those  of  the  univalent  metals,  having  the  form 
K2O,  if  we  leave  cuprous  oxide  and  aurous  oxide  out  of  account, 
have  the  most  strongly  basic  qualities.  Those  of  the  bivalent 
metals  of  the  form  MgO,  when  this  is  the  only  oxide  which  they 
furnish,  are  base-forming.  Those  of  the  trivalent  metals  of  the 
form  A1203,  known  as  sesquioxides  (Lat.  sesqui-,  one-half  more), 
are  the  least  basic  of  the  basic  oxides.  The  oxides  of  the  forms 
Sn02,  Sb20s,  CrOa,  and  Mn2C>7,  in  which  the  metals  have  valences 
from  4  to  7,  are  mainly  acid-forming  oxides,  although  the  same  ele- 
ments usually  have  other,  lower  oxides,  which  are  basic. 

The  hydroxides  are  formed,  in  the  cases  of  the  metals  of  the 
alkalies  and  alkaline  earths  by  direct  union  of  water  with  the  oxides. 
They  are  produced  also  by  double  decomposition  when  a  soluble 
hydroxide  acts  upon  a  salt  (cf.  p.  398).  All  hydroxides,  except 
those  of  the  alkali  metals,  lose  the  elements  of  water  when  heated, 
and  the  oxides  remain.  In  some  cases  the  loss  takes  place  by  stages, 
just  as  was  the  case  with  orthophosphoric  acid  (p.  558).  Thus  lead 
hydroxide  Pb(OH)2  (q.v.)  first  gives  the  hydroxide  Pb2O(OH)2,  then 
Pb3O2(OH)2,  and  then  the  oxide  PbO.  The  hydroxides  of  mercury 
and  silver,  if  they  are  formed  at  all,  are  evidently  unstable,  for, 
when  either  material  is  dried,  to  remove  adhering  moisture,  it  is 
found  to  contain  nothing  but  the  corresponding  oxide.  With  the 
exception  of  those  of  the  metals  of  the  alkalies  and  alkaline  earths, 
all  the  hydroxides  are  little  soluble  in  water. 

Compounds  of  the  Metals:  Salts.  —  It  may  be  said,  in  gen- 
eral, that  each  metal  may  form  a  salt  by  combination  with  each  one  of 
the  acid  radicals.  In  the  succeeding  chapters  we  shall  describe  only 
those  salts  which  are  manufactured  commercially,  or  are  of  special 
interest  for  some  other  reason.  The  various  salts  will  be  described 
under  each  metal.  Here,  however,  a  few  remarks  may  be  made 
about  the  characteristics  of  the  more  common  groups  of  salts.  The 
salts  are  classified  according  to  the  acid  radicals  which  they  con- 
tain. 


THE  BASE-FORMING  ELEMENTS  655 

The  chlorides  may  be  made  by  the  direct  union  of  chlorine  with 
the  metal  (cf.  p.  221),  or  by  the  combined  action  of  carbon  and  chlorine 
upon  the  oxide  (cf.  p.  632).  The  general  methods  for  making  any 
salt  (p.  213),  such  as  the  interaction  of  a  metal  with  an  acid,  or  of  the 
oxide,  hydroxide,  or  another  salt  with  an  acid,  or  the  double  decom- 
position of  two  salts,  may  be  used  also  for  making  chlorides.  The 
chlorides  are  for  the  most  part  soluble  in  water.  Silver  chloride, 
mercurous  chloride,  and  cuprous  chloride  are  almost  insoluble, 
however,  and  lead  chloride  is  not  very  soluble.  Most  of  the  chlorides 
of  metals  dissolve  without  decomposition,  but  hydrolysis  is  noticeable 
in  the  case  of  the  chlorides  of  the  trivalent  metals,  such  as  aluminium 
chloride  and  ferric  chloride  (cf.  p.  648).  The  chlorides  of  some  of  the 
bivalent  metals  are  hydrolyzed  also,  but,  as  a  rule,  only  when  they 
are  heated  with  water.  This  is  the  case  with  the  chlorides  of  mag- 
nesium, calcium,  and  zinc.  Most  of  the  chlorides  are  stable  when 
heated,  but  those  of  the  noble  metals,  particularly  gold  and  platinum, 
are  decomposed,  and  chlorine  escapes.  The  chlorides  are  usually 
the  most  volatile  of  the  salts  of  a  given  metal,  and  so  are  preferred  for 
the  production  of  the  spectrum  (q.v .)  of  the  metal,  and  for  fixing  the 
atomic  weight  of  the  metal  by  use  of  the  vapor  density.  Some  of  the 
metals  form  two  or  more  different  chlorides.  For  example,  indium 
gives  InCl,  InCl2,  and  InCl3. 

The  sulphides  are  formed  by  the  direct  union  of  the  metal  with  sul- 
phur, or  by  the  action  of  hydrogen  sulphide  or  of  some  soluble  sul- 
phide upon  a  solution  of  a  salt  (cf.  p.  421).  In  one  or  two  cases  they 
are  made  by  the  reduction  of  the  sulphate  with  carbon.  The  sul- 
phides, except  those  of  the  alkali  metals,  are  but  little  soluble  in  water. 
The  sulphides  of  aluminium  and  chromium  are  hydrolyzed  completely 
by  water,  giving  the  insoluble  hydroxides,  and  those  of  the  metals  of 
the  alkaline  earths  are  partially  hydrolyzed  (cf.  p.  421). 

Some  of  the  metals,  when  they  are  in  the  molten  form,  dissolve 
carbon,  and,  when  they  are  cooled  once  more,  deposit  it  in  the  form 
of  graphite.  This  is  true  particularly  of  platinum  and  iron.  The 
carbides  are  usually  formed  in  the  electric  furnace  by  interaction  of 
an  oxide  with  carbon  (cf.  p.  570).  Some  of  them  are  decomposed  by 
contact  with  water,  after  the  manner  of  calcium  carbide,  giving  a 
hydroxide  and  a  hydrocarbon.  Of  this  class  are  lithium  carbide  Li2C2, 
barium  and  strontium  carbides  BaC2  and  SrC*  aluminium  carbide 
AUCs,  manganese  carbide  MnC,  and  the  carbides  of  potassium  and 
glucinum.  Others,  such  as  those  of  molybdenum  MojC  and  chro- 
mium Cr3C2,  are  not  affected  by  water. 


656  INORGANIC  CHEMISTRY 

The  nitrates  may  be  made  by  any  of  the  methods  used  for  prepar- 
ing salts.  The  normal  nitrates  are  all  at  least  fairly  soluble  in  water. 

The  sulphates  are  made  by  the  methods  used  for  making  salts, 
and  in  some  cases  by  the  oxidation  of  sulphides.  They  are  all 
soluble  in  water,  with  the  exception  of  those  of  lead,  barium,  and 
strontium.  Calcium  sulphate  is  meagerly  soluble. 

The  carbonates  are  prepared  by  the  methods  used  for  making 
salts.  They  are  all  insoluble  in  water,  with  the  exception  of  those  of 
sodium  and  potassium.  The  hydroxides  of  aluminium  and  tin  are  so 
feebly  basic  and  so  insoluble  that  these  metals  do  not  form  carbonates 
which  are  stable  in  contact  with  moisture  (cf.  p.  648). 

The  phosphates  and  silicates  are  prepared  by  the  methods  used  in 
making  salts.  The  former  are  obtained  also  by  special  processes 
already  described  (p.  559).  With  the  exception  of  the  salts  of 
sodium,  ammonium,  and  potassium,  all  the  salts  of  both  these  classes 
are  insoluble. 

Solubility  of  Bases  and  Salts.  —  The  solubilities  of  a  few  salts 
at  various  temperatures  have  already  been  given  (p.  191).  For  the 
exact  solubilities  of  a  large  number  of  bases  and  salts  (142)  at  18°, 
see  the  Table  inside  the  cover,  at  the  front  of  this  book.  Each  square 
contains  two  numbers  expressing  the  solubility  of  the  compound  whose 
cation  stands  at  the  head  of  the  column  and  whose  anion  is  indicated 
at  the  side.  The  solubility  is  that  of  the  hydrate  stable  at  18°, 
where  such  exists.  The  upper  number  in  each  case  gives  the  number 
of  grams  of  the  anhydrous  salt  held  in  solution  by  100  c.c.  of  water. 
The  lower  number  shows  the  number  of  moles  in  1 1.  of  the  saturated 
solution,  and  indicates  therefore  the  concentration  in  terms  of  a  molar 
solution  as  unity  —  the  molar  solubility.  In  the  cases  of  the  less 
soluble  compounds  the  values  are  not  exact,  but  they  will  serve  to 
show  roughly  the  relative  solubilities  when  several  substances  are 
compared.  The  numbers  for  small  solubilities  have  been  abbreviated. 
Thus,  0.064  =  0.0000004. 

The  following  are  the  solubilities  (number  of  grams  in  100  c.c. 
water)  at  18°  of  two  additional  insoluble  substances  and  of  three 
acid  salts. 

Mercurous  chloride,  0.0s2  Sodium  bicarbonate  9.6 

(molar  sol'ty,  0.04!)  Potassium  bicarbonate  26.1 

Mercuric  iodide,      0.044  Potassium  bisulphate  50.0 

(molar  sol'ty,  0.04)  Sodium  bisulphate  (0°)  50.0 


THE  BASE-FORMING  ELEMENTS 


657 


It  will  be  seen  that  some  compounds,  like  zinc  chloride  and 
barium  iodide,  are  exceedingly  soluble;  that  others,  like  potassium 
chloride  and  barium  chlorate,  are  of  medium  solubility;  that  still 
others,  like  calcium  hydroxide  and  calcium  sulphate,  are  sparingly 
soluble;  and,  finally,  that  some,  like  calcium  oxalate  CaC2O4  and 
barium  chromate,  are  almost  insoluble.  The  reader  should  note  the 
fact,  however,  that  the  differences  in  solubility  even  amongst  the 
insoluble  salts  are  as  great  as  amongst  the  soluble  ones. 

Solubilities  at  different  temperatures  are  shown  in  the  diagram, 
Fig.  79,  p.  191. 

Hydrated  Forms  of  Salts  Commonly  Used.  —  In  the  table 
given  below,  the  figures  refer  to  the  number  of  molecules  of  water  in 
the  hydrates  which  are  deposited  by  aqueous  solutions  of  the  salts 
in  the  neighborhood  of  18°.  The  letter  h  means  that  the  compound 
is  stable  when  heated,  the  letter  a  that  it  is  not  affected  by  the  air, 
the  letter  d  that  the  salt  is  deliquescent,  and  the  letter  e  that  the 
hydrate  loses  water  spontaneously  in  an  open  vessel,  i.e.,  is  efflorescent. 

Composition  of  Hydrates  of  Salts 


K 

Na 

Li 

Ag 

Ba 

Sr 

Ca 

Mg 

Zn 

Cd 

Cu 

Pb 

Cl 

Oh 

Oft 

9*1 

Oh 

9*1, 

fie 

fid 

fid 

Hrf 

9&f 

9*1 

Oh 

Br    

Oh 

Oh 

Oh 

Oh 

9*1, 

fid 

fid, 

fid 

9*1. 

4e 

4e. 

Oh 

I 

Oh 

Oh 

Oh 

Oh 

9*1 

fid 

Od 

Xrf 

Od 

On 

Oh 

NO3 

On 

On 

Od 

On 

On 

4e 

4d 

fid 

fid 

4d 

fid 

On 

C1O3 

On 

On 

Od 

On 

In 

5r/ 

9*1 

fid 

fid 

9*1 

fid 

In 

BrO3 

On 

On 

Od 

On 

In 

1a 

la 

fie 

fin 

*>*1 

fin 

la 

IO3             

On 

3p 

Od 

On 

1(7 

fie 

fie, 

4n 

9*i. 

On 

}n 

Oa 

C2H3O2  

Od 

3f 

9*1 

On 

1a 

%n 

9*1 

4d 

3a 

3d 

1a 

3a 

SO4      

Oh 

1(V 

Oh 

On 

Oh 

Oh 

9,a 

7e, 

7e, 

9%a 

5a 

Oh 

CrO4   

Oh 

\0e 

9a 

On 

Oh 

Oh 

1a 

7e. 

Oh 

C2O4 

In 

On 

On 

On 

In 

On 

In 

9*i 

9*i 

JVr 

ia 

On 

C03     

Ud 

We 

Oa 

Oe 

Oh 

Oh 

Oa 

3e 

Oa 

Oa 

Oa 

Isomorphism.  —  All  substances  which  crystallize  in  one  of  the 
forms  belonging  to  the  regular  system  (p.  172)  must  necessarily  have 
identical  crystalline1  shapes.  Thus,  crystallized  specimens  of  sodium 
chloride  and  of  lead  sulphide  (galena)  in  their  natural  shapes  are 
cubical.  The  forms  found  in  other  systems,  however,  are  capable 
of  assuming  an  infinite  diversity  of  shapes.  The  relative  lengthening 
or  shortening  in  one  direction,  shown  by  the  square  prismatic  and 
hexagonal  forms  (p.  172),  for  example,  makes  it  possible  for  each 


658  INORGANIC  CHEMISTRY 

separate  substance  to  adopt  proportions  which  are  more  or  less  dif- 
ferent from  those  of  every  other  substance.  Each  substance,  not- 
belonging  to  the  regular  system,  does,  in  fact,  crystallize  invariably 
in  forms  based  upon  its  own  fundamental  proportions,  and  differs 
therefore  in  its  angles  from  all  other  such  substances  in  a  way  that  is 
clearly  recognizable  by  refined  measurement. 

Now  it  is  found  that  substances  which  are  chemically  somewhat 
similar  (see  below)  frequently  crystallize  in  the  same  system  and  show 
proportions  which  are  almost,  although  not  quite,  identical.  Fur- 
ther, two  or  more  such  substances,  when  the  approach  to  identity  in 
angles  is  not  accidental,  can  take  part  in  the  construction  of  one  and 
the  same  crystal.  A  crystal  of  one  such  substance  placed  in  a  solu- 
tion of  the  other  will  continue  to  grow,  and  in  doing  so  will  follow  the 
pattern  already  set,  and  simply  increase  in  dimensions  by  accretion  of 
the  new  material.  When  a  solution  containing  two  such  substances 
deposits  crystals,  the  structures  are,  not  some  of  them  of  one  material 
and  some  of  the  other,  but  are  all  made  up  of  both  in  a  ratio  deter- 
mined by  the  relative  amounts  of  the  substances  in  the  solution.* 
Substances  related  in  these  two  ways,  that  is,  having,  when  separate, 
crystalline  forms  which  are  closely  alike  and  being  capable  of  forming 
homogeneous  crystals  containing  varying  proportions  of  the  two  in- 
gredients, are  called  isomorphous  substances  (Gk.  IO-QS,  equal;  f^op^, 
form).  Thus,  potassium  permanganate  KMn04  and  potassium 
perchlorate  KC1O4  crystallize  in  the  rhombic  system,  forming  crys- 
tals with  very  similar  angles  (Fig.  70,  p.  173),  and  when  a  solution 
containing  both  is  allowed  to  evaporate  there  is  formed  but  one 
set  of  crystals  made  up  of  both  substances.  Similarly,  potassium 
iodide  and  ammonium  iodide  crystallize  in  cubes  of  the  regular  system 
and,  since  all  cubes  are  alike,  necessarily  show  absolutely  identical 
angles.  In  addition  to  this,  however,  they  crystallize  together  from 
a  solution  containing  both  salts.  Other  pairs  from  amongst  sub- 
stances belonging  to  the  regular  system  would  not  do  this.  These  two 
salts  are  therefore  isomorphous.  Substances  which  thus  form  mixed 
crystals  have  approximately  equal  molecular  volumes. 

In  the  course  of  our  study  of  the  compounds  of  the  metals  we 
shall  have  occasion  to  note  many  examples  of  isomorphism.  Thus  the 
heptahydrates  of  the  sulphates  of  many  of  the  bivalent  metals, 
such  as  ZnSO4,7H2O,  NiSO4,7H2O,  MgSO4,7H2O,  etc.,  belong  to  the 

*  In  general,  two  substances  which  are  absolutely  unrelated  may  be  deposited 
simultaneously  from  mixed  aqueous  solutions,  but  some  of  the  crystals  are  purs 
specimens  of  one  substance,  and  the  rest  are  pure  specimens  of  the  other. 


THE  BASE-FORMING  ELEMENTS  659 

rhombic  system,  and  form  an  isomorphous  set  of  substances  known 
as  the  vitriols  (q.v.). 

The  alums  (q.v.)  also  constitute  an  important  set,  and  crystallize, 
separately  and  together,  in  the  regular  system.  Amongst  minerals, 
lead  sulphide  PbS  and  silver  sulphide  Ag2S  form  a  common  isomor- 
phous pair,  and  nearly  all  natural  specimens  of  galena  (q.v.)  contain 
at  least  a  little  silver  sulphide. 

The  chemical  significance  of  isomorphism  was  at  first  exaggerated. 
Thus  the  elements  magnesium  and  iron  are  not  especially  similar  in 
their  chemical  relations,  excepting  that  both  are  bivalent;  yet  they 
form  several  pairs  of  compounds  which,  like  the  sulphates  (above), 
are  isomorphous.  Still,  in  practical  chemical  work  a  knowledge  of 
the  relations  of  a  substance  in  respect  to  isomorphism  is  indispensable. 
It  enables  us  to  predict  the  probable  impurities  in  a  homogeneous- 
looking  material,  for  non-isomorphous  substances  would  have  given  a 
heterogeneous  mixture  with  it.  It  assists  us  in  separating  and  puri- 
fying chemical  substances,  for  non-isomorphous  substances  can  be  sep- 
arated by  recrystallization  from  water  (p.  481),  or  by  washing  with 
water  or  some  other  solvent  (p.  484),  at  a  temperature  at  which  the 
solubilities  of  the  substances  are  different  (see  Potassium  nitrate). 
Isomorphous  substances,  however,  can  be  separated  only  by  conver- 
sion into  some  other  form  of  combination  in  which  the  property  is 
lacking.  Thus,  silver  sulphide  cannot  be  separated  from  lead  sulphide 
by  Pattinson's  process  (q.v.),  and  so  the  mixed  metals,  to  the  separa- 
tion of  which  the  process  is  applicable,  must  first  be  secured  by 
reduction. 

Some  chemists  regard  isomorphous  mixtures  as  solid  solutions. 

Exercises.  —  1.  Compare  the  electrical  conductivities  of  normal 
sodium  hydroxide  and  normal  acetic  acid  with  the  conductivity  of 
copper.  What  length  of  copper  wire  will  present  the  same  resistance 
as  1  cm.  of  each  of  these  solutions  when  the  cross-sections  are 
alike? 

2.  What  do  we  mean  by  saying  that  an  oxide  is  strongly  or  feebly 
basic,  or  that  it  is  acidic  (p.  654)? 

3.  What  is  meant  by  the  same  terms  when  applied  to  an  hy- 
droxide? 

4.  Compare  the  molar  solubilities  at  18°,  (a)  of  the  halides  of  silver 
and  (b)  of  the  carbonates  and  (c)  oxalates  of  the  metals  of  the  alkaline 
earths,  noting  the  relation  between  solubility  and  atomic  weight. 

5.  What  are  the  molar  concentrations  of  chloride-ion  (cf.  p.  183)  in 


660  INORGANIC  CHEMISTRY 

saturated  solutions  of  silver  chloride  and  lead  chloride  at  18°,  assum- 
ing complete  ionization  in  these  very  dilute  solutions? 

6.  How  does  the  behavior  of  complex  acids,  like  chloroplatinic 
acid  H^PtCle,  differ  from  that  of  acid  salts? 

7.  Formulate  (p.  399)  (a)  the  hydrolysis  of  ammonium  ortho- 
silicate  (made  by  mixing  solutions  of  ammonium  chloride  and  sodium 
orthosilicate) ;    (6)  the  hydrolysis  of  aluminium  carbonate,  made  by 
mixing  solutions  of  aluminium  sulphate  and  sodium  carbonate. 


CHAPTER  XXXIII 

METALLIC  ELEMENTS   OF  THE  ALKALIES:    POTASSIUM  AND 

AMMONIUM 

The  Metals  of  the  Alkalies.  —  The  metals  of  this  family  form 
an  homogeneous  group,  and  there  is  a  very  general  similarity  between 
the  properties  of  the  corresponding  compounds.  Some  of  the  physi- 
cal properties  of  the  elements  themselves  can  best  be  shown  in  tabular 
form. 


AT.  WT. 

SP.  GB. 

M.-P. 

B.-P. 

Lithium  Li      

6  94 

0  53 

186° 

above  red  heat 

Sodium  Na      

23.00 

0.97 

95  6° 

742° 

Potassium  K  

39.10 

0.86 

62  5° 

720° 

Rubidium  Rb     

85.45 

1.53 

38  5° 

Caesium  Cs    

132.81 

1.87 

26.5° 

270° 

It  will  be  seen  that  the  specific  gravities  of  the  elements  increase  with 
rising  atomic  weight,  while  the  melting-points  and  boiling-points  fall 
(cf.  p.  462).  A  table  including  all  the  physical  properties,  both  of  the 
elements  and  their  compounds,  would  show  similar  characteristics  in  a 
general  way,  with  here  and  there  noticeable  irregularities  such  as  that 
shown  by  the  specific  gravity  of  sodium.  For  example,  the  melting- 
points  of  the  hydroxides  are:  sodium  318.4°,  potassium  360.4°,  rubid- 
ium 301°,  caesium  272.3°.  Sodium  (m.-p.  95.6°)  and  potassium 
(m.-p.  62.5°)  give  together  an  alloy  which  is  liquid  at  room  tempera- 
ture (p.  644). 

The  Chemical  Relations  of  the  Metallic  Elements  of  the 
Alkalies.  —  The  metals  which  are  chemically  most  active  are  in- 
cluded in  this  group,  and  the  activity  increases  with  rising  atomic 
weight,  caesium  being  the  most  active  positive  element  of  all.  A 
freshly  cut  surface  of  any  of  these  metals  tarnishes  by  oxidation  as 
soon  as  it  is  exposed  to  the  air.  Indeed,  there  is  scarcely  time  to  see 
the  metallic  appearance  in  the  case  of  potassium  and  the  metals  fol- 
lowing it.  All  of  these  metals  decompose  water  violently  (cf.  p.  115), 

661 


662  INORGANIC  CHEMISTRY 

liberating  hydrogen.  The  hydroxides  which  are  formed  by  this  action 
are  exceedingly  active  bases,  that  is  to  say,  they  give  a  relatively  large 
concentration  of  hydroxide-ion  in  solutions  of  a  given  molecular  con- 
centration (p.  374).  Lithium  hydroxide  is  the  least  active.  In  the 
dry  form  these  hydroxides  are  not  decomposed  by  heating,  while  the 
hydroxides  of  all  other  metals  lose  water  more  or  less  easily.  All 
these  metals  seem  to  combine  with  hydrogen,  lithium  giving  the  most 
stable  compound.  The  hydrides,  however,  unlike  those  of  many  of 
the  non-metals,  are  not  ionogens,  and  consequently  do  not  give  acids 
when  dissolved  in  water.  In  all  their  compounds  the  metals  of  the 
alkalies  are  univalent. 

The  family  may  be  subdivided  into  two  minor  groups.  The  com- 
pounds of  potassium,  rubidium,  and  caesium  resemble  one  another 
closely,  while  those  of  sodium  and  lithium  are  sometimes  largely  diver- 
gent in  physical  properties.  Thus,  the  chlorides  of  the  potassium  set, 
not  only  crystallize  in  cubes,  but  can  form  mixed  crystals  with  one 
another  in  all  proportions.  They  are  isomorphous  (cf.  p.  658).  The 
same  is  true  of  the  bromides  and  of  the  iodides.  Sodium  chloride, 
although  crystallizing  in  cubes  likewise,  does  not  fqrni  mixed  crystals 
with  the  chlorides  of  the  potassium  set.  In  the  case  of  lithium, 
the  hydroxide  is  not  nearly  so  soluble  as  are  the  hydroxides  of  the 
other  metals,  and  the  metal  gives  also  an  insoluble  carbonate  and 
phosphate,  in  which  respect  it  resembles  magnesium  and  differs  from 
all  the  other  members  of  the  present  group. 

The  compounds  of  ammonium  will  be  discussed  in  connection  with 
those  of  potassium,  to  which  they  present  the  greatest  resemblance. 

The  solubilities  are  often  decisive  factors  in  connection  with  the 
preparation  and  use  of  salts.  The  reader  will  find  most  of  these  in  the 
Table  inside  the  front  cover,  or  in  the  diagram  on  p.  191,  and,  as  a 
rule,  the  values  will  not  be  repeated  in  the  descriptive  paragraphs. 

POTASSIUM  K 

. 

Occurrence.  —  Silicates  containing  potassium,  such  as  felspar 
and  mica  (p.  636),  are  constant  constituents  of  volcanic  rocks,  and 
from  the  weathering  of  these  rocks,  and  of  the  detritus  formed  from 
them  which  constitutes  a  large  part  of  the  soil,  the  potassium  used  by 
plants  is  obtained.  These  minerals  are  not  used  commercially  as 
sources  of  potassium  compounds.  The  salt  deposits  (see  below)  con- 
tain potassium  chloride,  alone  (sylvite)  and  in  combination  with  other 
salts,  and  most  of  the  compounds  of  potassium  are  manufactured 


POTASSIUM  AND  AMMONIUM  663 

from  this  material.  Part  of  our  potassium  nitrate,  however,  is  puri- 
fied Bengal  saltpeter  (p.  525).  Potassium  sulphate  occurs  also  in  the 
salt  layers,  and  is  used  directly  as  a  fertilizer. 

Preparation.  —  Potassium  was  first  made  by  Davy  (1807)  by 
bringing  the  wires  from  a  battery  in  contact  with  a  piece  of  moist 
potassium  hydroxide.  Globules  of  the  metal  appeared  at  the  negative 
wire.  This  process  has  again  come  into  use,  commercially,  molten 
potassium  chloride  being  the  substance  decomposed.  The  process 
which,  during  the  intervening  years,  furnished  potassium,  was  the 
heating  of  potassium  carbonate  with  finely  divided  carbon  in  small 
retorts : 

K2C03+2C->2K+3CO. 

The  vapor  of  the  metal  tends  to  combine  with  the  carbon  monoxide, 
forming  an  explosive  compound  KeCeOe,  and  the  yield  is  thus  reduced. 

Physical  Properties.  —  Potassium  is  a  silver-white  metal  which 
melts  at  62.5°.  It  boils  at  720°,  giving  a  greenish  vapor.  The  metal 
and  its  compounds  confer  a  violet  tint  upon  the  Bunsen  flame,  and  the 
spectrum  (q.v.)  shows  characteristic  lines. 

Chemical  Properties.  —  The  density  of  the  vapor  shows  the 
molecular  weight  of  potassium  to  be  about  40,  so  that  the  vapor  is  a 
monatomic  gas.  The  element  unites  violently  with  the  halogens, 
sulphur,  and  oxygen.  In  consequence  of  the  latter  fact  it  is  usually 
kept  under  petroleum,  an  oil  which  neither  contains  oxygen  itself, 
nor  dissolves  a  sufficient  amount  of  moisture  from  the  air  (c/.  p.  87) 
to  permit  much  oxidation  of  the  potassium  to  take  place. 

The  Hydride  KH.  —  When  hydrogen  is  passed  over  potassium 
heated  to  360°,  a  hydride  is  formed.  By  washing  the  solid  product 
with  liquefied,  dry  ammonia  the  excess  of  potassium  is  removed  and 
white  crystals  remain.  These  have  the  composition  KH.  On  ac- 
count of  the  ease  with  which  it  decomposes,  the  substance  behaves 
much  like  potassium  itself.  When  thrown  into  water,  for  example, 
it  gives  potassium  hydroxide,  and  the  hydrogen  is  liberated  (c/.  p.  703). 

Potassium  Chloride  KCl.  —  Sea-water  and  the  waters  of  salt 
lakes  contain  a  relatively  small  proportion  of  potassium  compounds. 
During  the  evaporation  of  such  waters,  however,  the  potassium  com- 


664  INORGANIC  CHEMISTRY 

pounds  tend  to  accumulate  in  the  mother-liquor  while  sodium  chloride 
is  being  deposited.  Thus,  when  the  Salt  Lake  in  Utah  shall  have 
finally  dried  up,  the  upper  (last  to  be  formed)  part  of  the  bed  of  salts 
which  it  will  leave  behind  will  contain  layers  rich  in  compounds  of 
potassium.  This  condition  is  realized  in  geological  deposits  which 
have  been  formed  in  the  same  way.  Thus,  at  Stassf urt,  near  Magde- 
burg, there  is  a  thickness  of  more  than  a  thousand  meters  of  common 
salt,  more  or  less  mixed  with  and  intersected  by  layers  of  sedimentary 
deposits.  Above  this,  and  therefore  deposited  last,  are  25-30  meters 
of  salt  layers  in  which  the  potassium  salts  are  chiefly  found,  while  over 
all  are  several  hundred  meters  of  sandstone.  Formerly  the  upper 
layers  were  simply  stripped  off  and  rejected.  Now,  however,  the 
revenue  obtained  from  the  products  of  these  layers  is  many  times 
greater  than  that  coming  from  the  rock  salt  below. 

The  forms  in  which  the  potassium  chloride  is  found  in  the  salt 
beds  are  sylvite  KC1  and  carnallite  KCl,MgCl2,6H2O.  The  latter  is 
heated  with  a  small  amount  of  water,  or  with  a  mother-liquor  obtained 
from  a  previous  operation  and  containing  sodium  and  magnesium 
chlorides.  The  magnesium  sulphate  which  it  contains  as  an  impurity 
remains  undissolved.  From  the  clear  liquid,  when  it  cools,  potassium 
chloride  is  deposited  first  and  then  carnallite.  The  former  is  taken 
out  and  purified,  and  the  latter  goes  through  the  process  again.  This 
potassium  chloride  is  the  source  from  which  most  of  our  potassium 
hydroxide  and  potassium  carbonate,  as  well  as  salts  of  minor  commer- 
cial importance,  are  made.  It  is  a  white  substance  crystallizing  in 
cubes,  melting  at  about  750°,  and  slightly  volatile  at  high  tempera- 
tures. 

Recently,  the  giant  kelps  of  the  Pacific  coast  have  been  used  as 
a  source  of  potassium  chloride.  The  dried  seaweed  contains  9  per 
cent  of  this  salt  and  about  0.1  per  cent  of  iodine. 

Potassium  Iodide  KI.  —  When  iodine  is  heated  in  a  solution  of 
potassium  hydroxide,  the  iodate  and  iodide  are  both  formed  (p.  487) : 

6KOH  +  3I2  ->  5KI  +  KI03  +  3H2O. 

The  dry  residue  from  evaporation  is  heated  with  powdered  carbon  to 
reduce  the  iodate,  and  all  the  iodide  can  then  be  purified  by  recrystal- 
lization.  Another  method  of  preparation  consists  in  rubbing  together 
iodine  and  iron  filings  under  water.  The  soluble  ferrous  iodide  Fel2 
thus  formed  is  then  treated  with  additional  iodine  and  gives  a  sub- 
stance Fe3Is,  intermediate  in  composition  between  ferrous  and  ferric 


POTASSIUM  AND  AMMONIUM  665 

iodides.  This  is  also  soluble.  When  potassium  carbonate  is  added 
to  the  solution,  a  hydrated  magnetic  oxide  of  iron  is  precipitated,  car- 
bon dioxide'  escapes,  and  evaporation  gives  potassium  iodide : 

Fe3I8  +  4K2CO3  +  4H20-»  SKI  +  Fe3(OH)8  +  4C02. 

The  salt  forms  large,  somewhat  opaque  cubes  (m.-p.  623°).  It  is 
used  in  medicine  and  for  precipitating  silver  iodide  Agl  in  photog- 
raphy. In  the  laboratory  it  is  used  whenever  an  iodide  is  required, 
for  example,  when  experiments  with  iodide-ion  are  to  be  made. 

The  aqueous  solution  takes  up  free  iodine,  forming  KI3,  in  equi- 
librium with  dissolved  iodine:  I3~  <=±  I~  +  I2  (dslvd).  The  mixture 
is  used  in  testing  for  starch,  and  in  reactions  in  which  a  solution  of  free 
iodine  is  required. 

The  Bromide  and  Fluorides.  —  Potassium  bromide  KBr  may  be 
made  in  either  of  the  ways  used  for  the  iodide.  It  crystallizes  in 
cubes.  It  is  used  in  medicine  and  for  precipitating  silver  bromide  in 
making  photographic  plates  (q.v.).  In  the  laboratory  it  is  always 
employed  when  a  bromide  is  needed  as  a  source  of  bromide-ion. 

The  fluoride  of  potassium  K2F2  may  be  obtained  by  treating  the 
carbonate  or  hydroxide  with  hydrofluoric  acid.  It  is  a  deliquescent, 
white  salt.  When  treated  with  an  equi-molecular  quantity  of  hydro- 
fluoric acid  it  forms  potassium-hydrogen  fluoride  KHF2,  a  white  salt 
which  is  also  very  soluble.  This  acid  salt  is  used  in  the  preparation 
of  pure  hydrofluoric  acid,  since  the  latter  is  liberated  from  it  as  a  vapor 
at  a  high  temperature. 

Potassium  chloride  is  the  least  soluble  of  the  halides  of  potassium, 
the  bromide,  fluoride,  and  iodide  coming  next  in  that  order.  The 
position  of  the  fluoride  as  the  third  in  order,  when  we  should  expect 
it  to  be  the  least  soluble  (p.  284),  shows  that  this  compound  is  some- 
what exceptional.  It  is  also  slightly  hydrolyzed  by  water,  as  if  it  were 
a  salt  of  a  dibasic  acid  (cf.  p.  439).  These  facts,  together  with  the 
existence  of  the  acid  fluoride,  lead  us  to  assign  to  it  the  formula  K2F2. 
Other  acid  fluorides  of  the  formulae  KH2F3  and  KH3F4  have  likewise 
been  made.  Since  potassium,  hydrogen,  and  fluorine  are  always  uni- 
valent,  and  no  ordinary  valence  is  thus  available  for  holding  together 
groupings  more  complex  than  KF  and  HF,  we  may  regard  all  these 
four  fluorides  of  potassium  as  molecular  compounds  (p.  530). 

Potassium  Hydroxide  KOH.  —  This  compound,  known  also  as 
caustic  potash  and  colloquially  as  potassium  hydrate  (p.  150),  was 


666  INORGANIC  CHEMISTRY 

formerly  made  entirely  by  boiling  potassium  carbonate  with  cal- 
cium hydroxide  suspended  in  water  (milk  of  lime) : 

Ca(OH)2  (solid)  **  Ca(OH)2  (dslvd)  *±  2OH~  +  Ca++)          m       ' 

K2C03  *±  2K+  +  C03=  P  CaC°3  (dslvd)  *  P*^ 

(solid). 

The  operation  is  conducted  in  iron  vessels,  because  porcelain,  being 
composed  of  silicates,  interacts  with  solutions  of  bases.  The  action  is, 
in  theory,  precisely  similar  to  that  of  sulphuric  acid  upon  barium  diox- 
ide (cf.  p.  317).  The  potassium  carbonate  corresponds  to  the  acid, 
being  completely  dissolved  from  the  beginning,  and  the  calcium  hy- 
droxide to  the  dioxide,  since  its  relative  insolubility  (0.17  g.  in  100  g. 
Aq)  enables  the  water  to  take  up  fresh  portions  into  solution  only 
when  the  part  dissolved  has  already  undergone  chemical  change. 
The  calcium  carbonate  which  is  precipitated  is  still  more  insoluble 
(0.0013  g.  in  100  g.  Aq)  than  the  hydroxide,  and  hence  the  action 
goes  forward.  The  action  as  a  whole  is  reversible,  for  a  reason  which 
will  be  explained  later  (p.  700),  and  consequently  such  an  amount  of 
water  is  employed  that  the  solution  at  no  time  contains  more  than 
about  ten  per  cent  of  potassium  hydroxide  (sp.  gr.  1.1).  The  con- 
clusion of  the  action  is  recognized  when  a  clear  sample  of  the  liquid  no 
longer  effervesces  on  addition  of  a  dilute  acid,  and  is  therefore  free 
from  potassium  carbonate.  After  the  precipitate  has  settled,  the 
potassium  hydroxide  is  obtained  by  evaporation  of  the  clear  liquid, 
K+  +  OH-  -»  KOH. 

Potassium  hydroxide  is  now  manufactured  by  electrolytic  processes. 
When  a  solution  of  potassium  chloride  is  electrolyzed,  chlorine  is 
liberated  at  the  anode,  and  hydrogen  and  potassium  hydroxide  at  the 
cathode.  The  necessity  of  keeping  those  two  sets  of  products  apart, 
since  by  their  interaction  potassium  hypochlorite  and  potassium 
chloride  would  be  formed  (cf.  p.  475),  has  made  the  devising  of  suitable 
apparatus  extremely  difficult.  In  one  type  of  apparatus  a  partition 
of  asbestos  cloth  divides  the  cell  into  two  parts.  In  some  cases  this  is 
placed  vertically,  and  in  others  horizontally.  In  the  latter  case  the 
anode  is  on  the  upper  side  of  the  partition,  in  order  that  the  chlorine 
as  it  is  liberated  may  ascend  to  the  surface  without  stirring  up  the 
liquid  or  having  occasion  to  pass  near  the  partition.  In  all  cases  the 
anode  is  made  of  graphite,  since  this  substance  is  less  easily  attacked 
by  chlorine  than  is  any  other,  and  the  cathode  is  made  of  iron,  a 
metal  which  best  resists  the  action  of  alkalies.  The  chlorine  is  used 
for  making  bleaching  powder.  Pure  brine  flows  in  continuously  at 
one  point,  and  a  solution  of  the  hydroxide  containing  much  undecom- 


POTASSIUM   AND  AMMONIUM 


667 


posed  chloride  flows  out  at  another.  The  Townsend-Baekeland, 
Briggs,  and  Du  Bois  cells  are  of  this  type. 

The  same  process  is  applied  to  sodium  chloride.  In  both  cases, 
the  chlorine  is  compressed  in  iron  cylinders  to  give  liquid  chlorine 
(cf.  p.  221)  or  is  used  to  prepare  hypochlorites  such  as  bleaching  pow- 
der. The  hydrogen,  in  some  cases,  is  also  utilized.  It  should  be 
noted  that  in  the  electrolysis  of  solutions  of  potassium  chloride  or 
sodium  chloride,  potassium  or  sodium  is  not  set  free.  The  salts 
are  highly  ionized,  and  their  ions  carry  the  current  efficiently  by 
migration  between  the  electrodes.  But,  when  potassium-ion  reaches 
the  cathode,  there  is  a  choice  between  discharging  either  the  K+,  or 
the  H+  of  the  water.  The  latter  is  much  more  easily  discharged 
(see  Chap.  XXXVIII),  and  so  hydrogen  is  liberated.  There  is  not 
nearly  so  much  hydrogen-ion  present  at  any  one  time  as  there  is 
potassium-ion,  but,  as  the  H+  is  discharged  and  removed,  the  instan- 
taneous ionization  of  more  water  restores  the  supply. 

The  Castner-Kellner  apparatus  (Fig.  140)  employs  a  different 
principle  very  ingeniously  for  the  separation  of  the  products.  The 
two  end  compartments  are  filled  with  brine  and  contain  the  graphite 


FIG.  140. 

anodes.  The  central  compartment  contains  potassium  (or  sodium) 
hydroxide  solution  and  the  iron  cathode.  The  positive  current  enters 
by  the  anodes,  and  the  chlorine  is  therefore  attracted  to  and  liberated 
upon  the  graphite:  2C1~  +  2©  — >C12T.  After  rising  through  the 
liquid  it  is  collected  for  the  manufacture  of  liquefied  chlorine  or  of 
bleaching  powder.  The  ions  of  potassium  or  of  sodium,  as  the  case 
may  be,  are  discharged  upon  a  layer  of  mercury  which  covers  the 
whole  floor  of  the  box,  and  the  free  metal  dissolves  in  the  mercury, 
forming  an  amalgam  (p.  644).  The  dilution  of  the  sodium  by  mer- 
cury permits  the  discharge  of  the  sodium-ion,  making  this  discharge 


668  INORGANIC  CHEMISTRY 

easier  than  that  of  hydrogen-ion.  The  layer  of  mercury  extends  be- 
neath the  partitions,  and  a  slight  rocking  motion  given  to  the  cell  by 
the  cam  causes  the  amalgam  to  flow  below  the  partition  into  the 
central  compartment.  Here  the  sodium  leaves  the  mercury  in  the 
form  of  sodium  ions  and  is  attracted  by  the  cathode.  Upon  this, 
hydrogen  from  the  water  (H+  +  OH~  +±  H2O)  is  discharged,  and  the 
residual  hydroxide-ion,  together  with  the  metal-ion,  constitutes  potas- 
sium or  sodium  hydroxide:  2K+  +  2OH~  +  2H+  +  20  -»  2K+  + 
2OH~  +  H2  t  .  A  slow  influx  of  salt  solution  at  one  point  and  over- 
flow of  the  alkaline  solution  in  the  central  cell  at  another,  is  main- 
tained. The  overflowing  liquid  contains  20  per  cent  of  alkali. 
Since  in  this  form  of  the  apparatus  there  is  no  undecomposed  chloride 
present  in  the  part  of  the  solution  which  contains  the  hydroxide, 
simple  evaporation  to  dryness  furnishes  the  solid  alkali.  The 
amount  of  mercury  required,  however,  involves  a  large  investment 
of  capital. 

Potassium  hydroxide  is  exceedingly  soluble  in  water,  and  conse- 
quently, instead  of  being  crystallized  from  solution,  the  molten  residue 
from  evaporation  is  cast  in  sticks.  When,  for  chemical  purposes,  the 
hydroxide  is  required  free  from  potassium  carbonate  and  other  impu- 
rities, it  is  dissolved  in  alcohol,  in  which  the  other  substances  are  not 
soluble.  Evaporation  of  this  solution  gives  pure  caustic  potash.  The 
hydroxide,  because  of  its  great  solubility  and  the  consequent  very  low 
vapor  tension  of  its  solution  (cf.  p.  197),  is  highly  deliquescent.  It 
also  absorbs  carbon  dioxide  from  the  air,  giving  potassium  carbonate. 
This  salt  is  itself  deliquescent,  and  consequently  a  syrupy  solution  of 
the  carbonate  is  the  final  result  of  weathering.  Solutions  of  the 
hydroxide  have  an  exceedingly  corrosive  action  upon  the  flesh,  de- 
composing it  into  a  slimy  mass  by  hydrolyzing  the  proteins,  which  are 
easily  decomposed  into  organic  acids.  In  solution,  the  base  is  highly 
ionized,  furnishing  a  high  concentration  of  hydroxide-ion.  Its 
aqueous  solution  is  therefore  used  with  salts  of  other  metals  for  pre- 
cipitation of  less  soluble  bases. 

The  Oxides.  —  The  simple  oxide  K2O  may  be  made  by  heating 
potassium  nitrate  or  nitrite  with  potassium  in  a  vessel  from  which  air 
is  excluded:  2KN03  +  10K  — »  6K2O  +  N2.  It  interacts  violently 
with  water,  giving  the  hydroxide.  When  exposed  to  the  air  it  unites 
spontaneously  with  oxygen,  and  a  yellow  peroxide  K2O4  is  formed. 

When  the  metal  burns  in  oxygen,  K2C>4,  a  yellow  solid  is  the  prod- 
uct. This  substance  interacts  violently  with  water,  giving  potassium 


POTASSIUM  AND  AMMONIUM  669 

hydroxide,  and  the  excess  of  oxygen  is  liberated.  With  perfectly  dry 
oxygen,  potassium  does  not  unite,  even  when  it  is  heated  strongly. 

Potassium  Chlorate  KCIOZ.  —  The  preparation  of  this  salt,  by 
interaction  of  potassium  chloride  with  calcium  chlorate,  has  already 
been  described  (p.  481).  It  is  also  made  by  electrolysis  of  potassium 
chloride  solution,  the  potassium  hydroxide  and  chlorine  which  are 
liberated  being  precisely  the  materials  required.  All  that  is  necessary- 
is  to  use  a  warm,  concentrated  solution  and  to  provide  for  the  mixing 
of  the  materials  generated  at  the  electrodes.  The  salt  crystallizes  out 
when  the  solution  cools. 

Potassium  chlorate  crystallizes  in  monoclinic  plates.  It  melts  at 
about  351°,  and  at  a  temperature  slightly  above  this  the  visible  libera- 
tion of  oxygen  begins.  Since  heat  is  given  out  by  the  decomposition, 
the  action  may  be  almost  explosive  if  large  amounts  of  the  material 
are  employed.  On  account  of  the  ease  with  which  its  oxygen  is 
liberated,  the  salt  is  employed  in  making  fireworks  and  as  a  compo- 
nent, along  with  antimony  trisulphide,  of  the  heads  of  Swedish 
matches.  With  acids  it  is  used  as  an  oxidizing  agent  on  account  of 
the  chloric  acid  which  is  set  free  (p.  480).  It  is  also  employed  in 
medicine  as  a  gargle. 

Potassium  perchlorate  KC1O4,  formed  by  the  heating  of  the  chlo- 
rate (p.  483),  gives  white  crystals  belonging  to  the  rhombic  system. 
Compared  with  the  chlorate,  on  account  of  the  greater  difficulty^ 
in  liberating  its  oxygen  by  heat,  it  finds  little  practical  application. 

By  adding  chlorine-water  to  potassium  carbonate  solution,  a! 
mixture  of  the  chloride  and  potassium  hypochlorite  is  formed: 

HC1  +  HC10  +  K2C03  <=±  KC1  +  KC10  +  H20  +  C02. 

The  carbonic  acid,  however,  is  not  completely  displaced  by  the 
HC1O,  which  is  a  feeble  acid,  and  so  some  HC1O  remains.  Hence, 
the  solution  is  used,  under  the  name  eau  de  Javel  (often  misspelt 
Javelle),  in  the  household  for  removing  stains. 

The  Bromate  and  lodate.  —  These  are  the  most  familiar  salts 
of  their  respective  acids.  The  mode  of  their  preparation  has  already 
been  described  (p.  487).  Potassium  iodate  KIOs  may  be  made  also 
very  conveniently  by  melting  together  potassium  chlorate  and  potas- 
sium iodide  at  a  low  temperature.  The  iodate  is  much  less  soluble 
than  the  chloride,  and  the  mixture  may  be  separated  by  crystalliza- 
tion from  water. 


670  INORGANIC  CHEMISTRY 

Potassium  Nitrate  KNOz.  —  The  formation  of  this  salt  in 
nature  and  its  mode  of  extraction  and  purification  have  already  been 
described  (p.  525).  This  source  of  supply  proved  insufficient,  for 
the  first  time,  during  the  Crimean  war  (1852-55),  and  a  method  of 
manufacture  from  Chile  saltpeter  (sodium  nitrate),  which  is  a  much 
cheaper  substance,  was  introduced.  Sodium  nitrate  and  potassium 
chloride  are  heated  with  very  little  water,  and  the  sodium  chloride 
produced  by  the  action,  which  is  a  reversible  one,  is  by  far  the  least 
soluble  of  the  four  salts  (see  diagram,  p.  191).  On  the  other  hand, 
in  hot  water,  the  potassium  nitrate  is  by  far  the  most  soluble.  Hence 
the  hot  liquid  drained  from  the  crystals  contains  the  required  salt, 
and  much  of  the  sodium  chloride  is  in  the  form  of  a  precipitate.  If 
the  solubility  curve  of  potassium  nitrate  (p.  191)  is  examined,  it  will 
be  seen  that  this  salt  is  but  slightly  soluble  in  cold  water,  and  hence 
most  of  it  is  deposited  when  the  solution  cools.  The  crystals  are 
mixed  with  little  sodium  chloride,  for,  as  the  curve  shows,  common 
salt  is  little  less  soluble  at  10°  than  it  is  at  100°. 

Potassium  nitrate  gives  long  prisms  belonging  to  the  rhombic  sys- 
tem (Fig.  71,  p.  173).  It  melts  at  about  339°,  and  when  more  strongly 
heated  gives  off  oxygen,  leaving  potassium  nitrite  KNC>2  (p.  537). 
Although  it  does  not  form  a  hydrate,  the  crystals  enclose  small  por- 
tions of  the  mother-liquor,  and  consequently  contain  both  water  and 
impurities.  When  heated,  the  crystals  fly  to  pieces  explosively 
(decrepitate),  on  account  of  the  vaporization  of  this  water.  Many 
substances  which  form  large  crystals,  and  do  not  melt  when  warmed, 
behave  in  the  same  way  and  for  the  same  reason.  In  consequence  of 
this,  the  purest  salt  is  made  by  violent  stirring  of  the  solution  during 
the  operation  of  crystallization,  the  result  being  the  formation  of  a 
crystal-meal. 

Potassium  nitrate  is  used  chiefly  in  the  manufacture  of  gunpowder, 
which  contains  75  per  cent  of  the  highly  purified  salt.  The  other  com- 
ponents are  10  per  cent  of  sulphur,  14  per  cent  of  charcoal,  and  about 
1  per  cent  of  water.  The  ingredients  are  intimately  mixed  in  the 
form  of  paste,  and  the  material  when  dry  is  broken  up  and  sifted, 
grains  of  different  sizes  being  used  for  different  purposes.  The  chemi- 
cal action  which  takes  place  when  gunpowder  is  fired  in  an  open  space 
probably  results  chiefly  in  the  formation  of  potassium  sulphide,  car- 
bon dioxide,  and  nitrogen: 

2KN03  +  3C  +  S  -»  K2S  +  3C02  +  N2. 
The  explosion  occurring  in  firearms  follows  a  much  more  complex 


POTASSIUM  AND  AMMONIUM  671 

course,  and  half  of  the  solid  product  is  said  to  be  potassium  carbonate 
(a  solid,  hence  the  smoke).  One  gram  yields  264  c.c.  of  gases  (0°  and 
760  mm.),  and  a  much  larger  volume  at  the  temperature  of  the  ex- 
plosion, and  gives  660  calories.  The  pressure,  at  the  temperature  of 
the  explosion,  if  the  gases  could  be  confined  within  the  volume  origi- 
nally occupied  by  the  gunpowder,  would  reach  about  forty-four  tons 
per  square  inch.  In  recent  years  common  gunpowder  has  been  dis- 
placed largely  by  smokeless  powder  (pp.  528,  541),  which,  in  decom- 
posing, produces  no  solids. 

Potassium  nitrate  is  used  also  in  preserving  ham  and  corned 
beef,  on  which  it  confers  a  red  color. 

Paper  saturated  with  potassium  nitrate  solution  and  dried,  is 
known  as  touch-paper.  The  salt  interferes  with  the  access  of  air  to 
the  cellulose,  and  the  oxygen  for  the  combustion  is  obtained  from  the 
nitrate.  The  product  consequently  does  not  blaze,  yet  cannot  be 
blown  out. 

Potassium  Carbonate  K2COS.  —  This  salt  is  manufactured 
from  potassium  chloride,  which  is  heated  with  magnesium  carbonate 
(magnesite),  water,  and  carbon  dioxide  under  pressure: 

2KC1  +  3MgC03  +  C02  +  5H20  4  2KHMg(C03)2,4H20  J  +  MgCl2. 

The  hydrated  mixed  salt  is  separated  from  the  liquid  containing 
magnesium  chloride  and  decomposed  by  heating  with  water  at  120°. 
The  product  is  a  solution  of  potassium  carbonate,  from  which  the 
precipitated  magnesium  carbonate  is  removed  by  filtration  and  used 
again.  A  certain  amount  is  also  obtained  from  the  fatty  material, 
known  as  suint,  which  forms  about  50  per  cent  of  the  weight  of  sheep's 
wool.  The  suint  .is  separated  from  the  latter  by  washing.  When 
this  material,  which  contains  the  potassium  salt  of  sudoric  acid  in 
large  proportions,  is  calcined,  potassium  carbonate  remains,  and  is 
extracted  from  the  ash  with  water.  Some  plants,  like  the  sugar-beet, 
take  up  exceptional  quantities  of  potassium  salts  from  the  soil.  The 
molasses  remaining  from  the  crystallization  of  beet-sugar  (p.  606) 
is  mixed  with  yeast  and  fermented.  After  the  alcohol  has  been  dis- 
tilled off,  the  liquid,  containing  organic  salts  of  potassium  in  solution, 
is  evaporated,  and  the  residue  is  ignited.  In  some  districts  potassium 
carbonate  is  still  extracted  from  wood-ashes,  its  original  source  and 
the  origin  of  its  name,  potash. 

This  salt  is  usually  sold  in  the  form  of  an  anhydrous  powder  (m.-p. 
over  1000°) .  When  crystallized  from  water  it  gives  a  hydrate  2K2CO3) 


672  INORGANIC  CHEMISTRY 

3H2O.  It  is  extremely  deliquescent.  Its  aqueous  solution  has  an 
alkaline  reaction.  The  hydrolysis  of  the  salt  by  the  water  is  exactly 
analogous  to  that  of  sodium  sulphide  (p.  418),  although  much  slighter. 
The  more  elaborate  scheme  given  in  that  connection  may  be  put  in 
simpler  form  to  show  that  the  action  consists  essentially  in  the  for- 
mation of  the  ion  HCO3~,  by  union  of  the  ion  CO3=  with  the  H+  of  the 
water.  This  takes  place  because  the  ionization  of  the  HCO3~  ion  is 
small  enough  to  be  commensurable  with  that  of  water  itself:  CO3~  + 
H+  +  OH~  ->  HCOr  +  OH~.  The  commercial  name  of  the  sub- 
stance is  pearl  ash.  It  is  used  in  making  soft  soap  and  hard  (in- 
fusible) glass.  It  is  also  employed,  by  interaction  with  acids,  in 
making  salts  of  potassium. 

The  use  of  the  bicarbonate  KHC03  in  purifying  carbon  dioxide 
has  already  been  mentioned  (p.  573).  Before  the  nineteenth  century, 
this  salt  was  used  under  the  name  saleratus  (Lat.  aerated  salt),  a 
name  now  sometimes  given  to  the  baking  soda  NaHC03  which  has 
displaced  it. 

When  a  concentrated  solution  of  the  salt  is  electrolyzed  in  such  a 
way  that  the  anode,  towards  which  the  KCO3~  ions  travel,  consists  of 
a  thin  platinum  wire,  the  crowding  together  of  the  discharged  ma- 
terial results  in  the  formation  of  the  percarbonate  (cf.  p.  449)  : 

+  2e-»K2C206. 


The  operation  must  be  conducted  between  —  15°  and  0°.  When  the 
solution  in  the  porous  cell  surrounding  the  anode  is  evaporated,  the 
product  is  obtained  as  an  amorphous  bluish-white  powder.  The  sub- 
stance liberates  oxygen  when  heated,  and  in  other  respects  behaves 
like  the  persulphates.  When  it  is  treated  with  a  dilute  acid,  a  solu- 
tion containing  hydrogen  peroxide  is  formed.  The  compound  is 
therefore  a  mixed  anhydride  (p.  449)  of  hydrogen  peroxide  and  potas= 
sium  bicarbonate. 

Potassium  Cyanide.  —  Formerly  this  compound  was  made  by 
heating  potassium  carbonate  with  nitrogenous  animal  matter.  So 
many  other  substances  were  formed  at  the  same  time,  however,  that 
the  required  product,  which  is  very  soluble,  was  difficult  to  isolate  in 
a  state  of  purity.  It  is  now  made  by  heating  together  potassium 
ferrocyanide  (q.v.)  with  sodium:  K4Fe(CN)6  +  2Na  -»  4KNC  + 
2NaNC  +  Fe,  which  gives  a  mixture  of  both  cyanides.  When  the 
residue  is  extracted  with  water,  only  the  cyanides  dissolve,  and  are 
easily  crystallized  in  pure  form  from  the  solution.  Very  interesting 


POTASSIUM  AND  AMMONIUM  673 

ig  the  formation  of  potassium  cyanide  in  the  blast  furnace  (q.v.). 
Carbon  and  nitrogen  unite  at  a  very  high  temperature  to  form 
cyanogen  (p.  625),  which  is  an  endothermal  compound  (van't  HofFs 
law,  p.  305),  and  a  sufficient  amount  of  potassium  is  found  in  the 
materials  to  complete  the  production  of  the  salt. 

Potassium  cyanide  crystallizes  in  cubes.  It  is  extremely  soluble 
in  water,  and  is  therefore  deliquescent.  Its  poisonous  qualities  are 
equal  to  those  of  hydrocyanic  acid.  The  acid  is  so  feeble  as  to  be 
liberated  even  by  the  carbon  dioxide  of  the  air,  and  hence  this  salt 
always  has  a  distinct  odor  of  hydrocyanic  acid.  Potassium  cyanide 
has  a  great  tendency  to  form  complex  compounds  with  cyanides  of 
other  metals  (cf.  p.  649).  Complex  compounds  of  this  kind  were  used 
in  the  galvanic  deposition  of  silver  and  gold  in  commercial  electro- 
plating, and  were  formed  when  the  cyanide  was  used  in  extracting 
gold  (q.v.)  from  its  ores.  For  these  purposes,  sodium  cyanide,  which 
is  now  much  cheaper,  has  entirely  displaced  the  potassium  salt.  The 
tendency  to  form  complex  compounds  is  doubtless  connected  with 
the  fact  that  the  cyanides  are  unsaturated  compounds  in  which  the 
carbon  has  two  free  valences:  K  —  N  =  C  (p.  626). 

Potassium  cyanate  KNCO  is  made  by  heating  potassium  cyanide 
in  the  air,  or,  still  better,  with  some  easily  decomposed  oxide  (p.  626). 
It  is  a  white,  easily  soluble,  crystalline  salt. 

Potassium  thiocyanate  KNCS  may  be  obtained  by  melting  potas- 
sium cyanide  with  sulphur  (cf.  p.  626).  It  is  a  white,  deliquescent 
salt  which  finds  some  applications  in  chemical  analysis. 

The  Sulphate  and  Bisulphate.  —  The  sulphate  of  potassium 
is  a  constituent  of  several  double  salts  found  in  the  Stassfurt  deposits. 
It  is  extracted  from  schoenite  MgSO4,K2SO4,6H2O  and  kainite  MgSO4,- 
MgCl2,K2SO4,6H2O.  The  former  is  treated  with  potassium  chloride 
and  comparatively  little  water,  whereupon  the  relatively  insoluble 
potassium  sulphate  crystallizes  out,  and  the  magnesium  chloride 
remains  in  the  mother-liquor.  The  crystals  belong  to  the  rhombic 
system,  contain  no  water  of  crystallization,  and  melt  at  1066°.  This 
salt  is  employed  in  making  potassium  carbonate  by  the  Le  Blanc 
process  and  in  preparing  alum  (q.v.).  It  is  also  much  used  as  a  fer- 
tilizer. Since  plants  take  up  solutions  through  their  cell  walls,  they 
can  absorb  soluble  compounds  only.  They  are,  therefore,  dependent, 
for  the  potassium  compounds  which  they  require,  upon  the  weathering 
out  of  soluble  potassium  compounds  from  the  insoluble  potassium 
silicates  contained  in  the  soil.  '  The  weathering  takes  place  too  slowly 


674  INORGANIC   CHEMISTRY 

to  furnish  a  sufficient  supply  for  many  crops,  particularly  that  of  the 
sugar-beet.  Hence  potassium  sulphate  is  mixed  directly  with  the  soil. 
The  mineral  kainite  itself  is  used  for  the  same  purpose. 

Potassium-hydrogen  sulphate  (bisulphate)  KHSO4  is  made  by 
the  action  of  sulphuric  acid  upon  potassium  sulphate:  K2S04  + 
£[2864  — >  2KHSO4.  It  crystallizes  from  water,  in  which  it  is  very 
soluble,  in  tabular  crystals.  When  heated  to  about  200°  it  melts,  and 
the  elements  of  water  are  eliminated,  the  pyrosulphate  remaining: 
2KHS04  ->  H20  +  K2S2O7.  The  latter,  when  still  further  heated, 
yields  sulphur  trioxide  and  potassium  sulphate.  The  bisulphate  is 
used  in  analysis  for  the  purpose  of  interacting  with  oxides  and  silicates 
and  converting  them  into  sulphates.  The  substance  is  more  efficient 
than  sulphuric  acid  for  this  purpose,  because  the  latter  cannot  be 
heated  above  330°,  while  the  liberation  of  the  active  sulphur  trioxide 
from  this  salt  takes  place  at  a  bright-red  heat.  The  aqueous  solu- 
tion of  the  bisulphate  is  strongly  acid  on  account  of  the  considerable 
ionization  of  the  hydrosulphate-ion. 

Sulphides  of  Potassium.  —  By  the  treatment  of  a  solution  of 
potassium  hydroxide  with  excess  of  hydrogen  sulphide,  a  solution  of 
potassium-hydrogen  sulphide  is  obtained.  Evaporation  of  the  solu- 
tion gives  a  deliquescent  solid  hydrate  2KHS,H2O.  When  the  solu- 
tion, before  evaporation,  is  treated  with  an  equivalent  amount  of 
potassium  hydroxide,  and  the  water  is  driven  off,  the  sulphide  K2S 
remains  behind  (cf.  p.  418) : 

KHS  +  KOH  +±  K2S  +  H20  T . 

With  proper  care,  the  very  soluble  hydrate  K2S,5H20  may  be  ob- 
tained. Considerable  amounts  of  sulphur  can  be  dissolved  in  solu- 
tions of  either  of  these  sulphides.  By  evaporation  of  the  resulting 
yellow  liquids,  various  polysulphides  have  been  obtained.  To  some 
of  these  have  been  ascribed  the  formulae  K2S3,  K4S7,  K2S4,  K4Sg,  and 
K2Ss,  but  they  are  probably  K2Ss  (which  has  been  isolated),  or  mix- 
tures of  the  pentasulphide  with  K2S.  Similar  substances  are  pro- 
duced, as  a  result  of  the  liberation  and  recombination  of  sulphur, 
when  the  solutions  are  exposed  to  the  oxidizing  action  of  the  air: 

2KHS  +  02  ->  2KOH  +  2S. 

In  most  respects  the  corresponding  compounds  of  potassium  and 
sodium  are  similar  in  their,  physical  properties  and  chemical  action. 
Since,  however,  the  latter  are  almost  Uniformly  less  expensive,  they 


POTASSIUM   AND  AMMONIUM  675 

find  much  wider  application.  In  a  few  cases,  however,  the  potassium 
salt  is  more  generally  used.  Thus,  potassium  chlorate  and  potassium 
iodide  are  much  less  soluble  than  the  corresponding  sodium  com- 
pounds, and  it  is  consequently  possible  in  each  of  these  two  cases  to 
separate  by  crystallization,  and  to  purify  the  potassium  salt  with 
greater  ease. 

Properties  of  Potassium-ion  K+:   Analytical  Reactions.  — 

The  positive  ionic  material  of  "the  potassium  salts  is  a  colorless  sub- 
stance. It  unites  with  all  negative  ions,  and  most  of  the  resulting 
compounds  are  fairly  soluble.  For  its  recognition  we  add  solutions 
containing  those  ions  which  give  with  it  the  least  soluble  salts.  Thus, 
with  chloroplatinic  acid  E^PtCle  it  gives  a  yellow  precipitate  of  potas- 
sium chloroplatinate  K^PtCle.  Since  nearly  one  part  of  this  salt 
dissolves  in  100  parts  of  water,  the  test  is  far  from  being  a  delicate  one. 
The  solubility  in  alcohol  is  much  smaller,  and  consequently  the  pre- 
cipitate may  frequently  be  obtained  from  a  dilute  solution  by  adding 
more  than  an  equal  volume  of  alcohol.  Picric  acid  (p.  527)  gives 
potassium  picrate  KOCeH^NC^s,  which  is  much  less  soluble  in  water 
(0.4  parts  in  100  at  15°).  Perchloric  acid  HC1O4  and  hydrofluosilicic 
acid  H2SiFe  likewise  give  somewhat  insoluble  salts  of  potassium. 
Potassium-hydrogen  tartrate  KHC4H4Oe  is  precipitated  by  the  addi- 
tion of  tartaric  acid  to  a  sufficiently  concentrated  solution  of  a  potas- 
sium salt.  The  normal  tartrate  K^H^Oe  is  much  more  soluble.  It 
may  be  obtained  by  treating  the  precipitate  with  a  solution  of  potas- 
sium hydroxide.  Addition  of  an  acid  to  this  solution  causes  reprecipi- 
tation  of  the  bitartrate. 

The  Spectroscope.  —  A  much  more  delicate  test  for  the  recog- 
nition of  a  potassium  compound  consists  in  the  examination  by  means 
of  the  spectroscope  of  the  light  given  out  by  a  Bunsen  flame,  in 
which  a  little  of  the  salt  is  held  upon  a  platinum  wire.  When  the 
amount  of  potassium  is  considerable,  and  no  other  substance  which 
would  likewise  color  the  flame  is  present  to  mask  the  effect,  the 
violet  tint  is  recognizable  by  the  eye.  In  general,  however,  the  light 
must  be  analyzed. 

White  light  is  composed  of  vibrations  of  every  wave-length  within 
a  certain  range.  If  the  light  is  made  up  of  a  few  wave-lengths  only, 
it  appears  to  the  eye  to  be  colored.  Now,  when  a  narrow  bundle 
of  rays  of  white  light,  coming  through  a  slit,  falls  upon  a  three-sided 
prism  standing  with  its  edges  parallel  to  the  slit,  the  rays  of  various 


676  INORGANIC  CHEMISTRY 

wave-lengths  are  retarded  to  different  extents  as  they  pass  through  the 
glass,  and  in  consequence  are  bent  from  their  paths  by  varying 
amounts.  Fig.  141  shows  a  horizontal  section  through  the  slit  (S)  and 
prism,  in  which  the  width  of  the  slit  and  of  the  beam  of  light  are 

exaggerated.  The  light  emerging  at 
the  other  side  of  the  prism  consists, 
therefore,  of  a  series  of  images  of  the 
slit  arranged  side  by  side.  The  red 
light  is  least  refracted,  and  the  red 
images  of  the  slit,  therefore,  are  most 
nearly  in  the  same  straight  line  with 
the  original  beam.  The  yellow,  green, 
blue,  and  violet  images  are  displaced 
FIG  141  more  and  more  from  this  direction,  and 

the  resulting  colored  band  is  called  a 

spectrum.  The  whole  series  of  images  of  the  slit  may  be  received 
upon  a  screen,  or  directly  upon  an  eye  looking  towards  the  prism. 
Now,  when  the  light  comes  from  the  vapor  of  potassium  heated  in  a 
Bunsen  flame,  there  are  produced,  not  thousands  of  images  of  the  slit, 
representing  as  many  different  wave-lengths  of  light,  but  only  two 
images,  one  red,  and  one  deep  blue,  corresponding  to  the  two  wave- 
lengths which  are  alone  contained  in  the  original  light.  In  a  more 
powerful  instrument  other  fainter  lines  are  seen  also.  Naturally  the 
brightness  of  all  these  lines  is  together  equal  to  that  of  the  original 
beam.  No  other  substance  gives  any  of  those  particular  lines,  al- 
though many  others  give  blue  and  red  light  of  somewhat  different 
wave-lengths.  Thus,  strontium  compounds  give  a  blue  light  along 
with  several  red  tints,  but  when  strontium  and  potassium  are  used 
together,  the  lines  are  found  not  to  be  coincident.  In  the  case  of 
strontium,  all  the  lines  lie  nearer  to  the  yellow  than  in  that  of  potas- 
sium. Since  the  whole  light  of  the  compound  is  thus  concentrated 
in  one  or  two  narrow  strips,  easily  visible  against  a  dark  background, 
small  amounts  of  the  elements  give  effects  which  are  readily  recogniz- 
able in  the  instrument.  This  remains  true  even  when,  to  the  eye,  the 
colors  are  completely  obscured  by  the  much  more  brilliant,  yellow 
light  which  compounds  of  sodium  produce.  In  the  spectrum  of 
sodium,  this  yellow  light  is  all  concentrated  into  two  yellow  lines 
which  lie  very  close  together. 

Helium  gives  many  lines,  but  one  orange  line  (D3),  in  particular, 
was  noted  in  the  spectrum  of  the  sun's  photosphere  many  years  before 
the  element  was  obtained  from  terrestrial  sources  by  Ramsay.  When 


POTASSIUM   AND  AMMONIUM  677 

the  spectra  of  helium  and  other  gases  are  to  be  examined  in  the 
laboratory,  a  little  of  the  material  is  enclosed  in  a  narrow,  exhausted 
tube,  through  which  an  electrical  discharge  can  be  passed  between 
platinum  wires.  Under  this  treatment  helium  shows  its  conspicuous 
orange  line,  and  hydrogen  a  red  and  two  blue  ones.  In  this  apparatus 
compounds  are  dissociated  and  give  the  spectra  of  their  constituents. 
When  a  Bunsen  flame  is  used  with  the  salts*  of  metals,  however,  the 
temperature  is  not  high  enough  to  render  visible  the  spectra  of  the 
non-metals  contained  in  them.  Indeed,  even  of  the  metals  them- 
selves, only  the  members  of  the  alkali  and  alkaline-earth  groups  give 
distinct  results. 

RUBIDIUM  AND  CAESIUM 

Soon  after  the  invention  of  the  spectroscope  by  Bunsen  and  Kirch- 
hoff,  the  instrument  was  applied  to  the  examination  of  many  sub- 
stances. In  1860  Bunsen  discovered  several  new  lines  in  the  spectrum 
given  by  materials  derived  from  the  salts  in  Durkheim  mineral  water. 
Two  new  elements  of  the  alkali  group  were  found  to  cause  their 
presence,  and  were  named,  from  the  colors  of  the  lines  which  they 
gave,  rubidium  (red)  and  caesium  (blue).  Both  elements  have  since 
been  found  in  small  quantities  in  various  minerals.  Rubidium  is 
obtainable  with  relative  ease  from  the  mother-liquors  of  the  Stassfurt 
works. 

The  metals  may  be  obtained  by  heating  their  hydroxides  with 
magnesium  powder.  The  salts  of  these  two  elements  are,  in  crystal- 
line form  and  solubility,  very  much  like  those  of  potassium.  In  some 
cases  the  difference  in  solubility  is  sufficient  to  make  separation  pos- 
sible. Thus,  a  mixture  containing  compounds  of  these  two  metals 
and  of  potassium  gives  with  chloroplatinic  acid  a  yellow  precipitate, 
consisting  of  the  three  insoluble  chloroplatinates.  The  solubilities  at 
10°,  however,  are  as  follows :  Potassium  chloroplatinate  0.9,  rubidium 
chtoroplatinate  0.15,  caesium  chloroplatinate  0.05.  Hence,  when  the 
mixed  precipitates  are  carefully  washed  with  small  quantities  of  cold 
water  the  potassium  chloroplatinate  can  be  almost  entirely  removed. 
On  similar  principles  the  two  other  metals  can  be  separated  from  one 
another.  The  iodides  of  all  three  elements  combine  with  iodine, 
giving  tri-iodides  (cf.  p.  665),  of  which  the  tri-iodide  of  caesium  is  the 
most  stable.  The  extra  iodine  must  be  held  to  have  entered  into 

*  The  chlorides  are  preferred  because  of  their  volatility.  The  salts  of  the  oxy- 
gen acids  are  dissociated,  and  leave  the  highly  involatile  oxides  (e.g.,  pp.  430,  531). 


678  INORGANIC  CHEMISTRY 

combination  with  the  iodine  of  the  compound,  and  not  with  the  metal. 
In  the  parallel  case  of  hydriodic  acid,  the  union  with  extra  iodine  (p. 
279)  seems  to  show  conclusively  that  iodide-ion  can  combine  with 
iodine.  While  an  inclination  to  trivalence  in  one  of  the  metals  of  the 
alkalies  would  furnish  a  very  acceptable  link  between  the  two  sides 
of  the  first  column  in  the  periodic  table  (p.  466),  since  gold  is  a  triva? 
lent  element,  no  such  tendency  has  been  proven. 

AMMONIUM 

The  compounds  of  ammonium  claim  a  place  with  those  of  the 
alkali  metals  because  in  aqueous  solution  they  give  ammonium-ion 
NH4+,  an  ion  which  in  its  behavior  closely  resembles  potassium-ion. 
Some  of  the  special  properties  peculiar  to  ammonium  compounds,  and 
particularly  the  properties  of  ammonium  hydroxide  NH4OH,  have 
been  discussed  in  detail  already  (pp.  519-521;. 

Ammonium  Chloride  NH^Cl.  —  This  salt,  known  commer- 
cially as  sal  ammoniac,  like  all  the  other  compounds  of  ammonium,  is 
prepared  from  the  ammonia  dissolved  by  the  water  used  to  wash 
illuminating-gas  (p.  613).  It  is  purified  by  sublimation,  and  then 
forms  a  compact  fibrous  mass.  It  crystallizes  from  solution  in  forms  of 
the  regular  system,  which  are  often  arranged  according  to  a  feathery 
pattern.  At  337.8°  its  vapor  exercises  one  atmosphere  pressure,  and 
is  dissociated  into  ammonia  and  hydrogen  chloride  to  the  extent  of  62 
per  cent  (p.  520). 

Ammonium  bromide  NH4Br  and  ammonium  iodide  NH4I  are 
white  salts  which  crystallize  in  the  regular  system,  and  are  isomor- 
phous  with  the  corresponding  potassium  salts.  They  are  dissociated 
by  heat,  but  the  degree  of  dissociation  of  the  bromide  diminishes  from 
320°  upwards  (p.  305)  and  the  vapor  of  the  iodide  contains  much 
(NHJJ2,  and  a  little  NH3  +  HI. 

Ammonium  Hydroxide  NH*OH.  —  The  nature  and  behavior 
of  this  substance  have  been  fully  discussed  (p.  520).  It  may  be  re- 
marked here  that  its  very  small  basic  activity  as  compared  with  that 
of  potassium  hydroxide  is  only  in  part  due  to  the  low  degree  of  ion- 
ization  of  its  molecules.  A  normal  solution  of  ammonia  contains 
much  free  NH3,  besides  the  NH4OH  produced  by  its  union  with  water. 
Measurement  shows  that  two-thirds  of  the  ammonia  is  not  actually 
in  the  form  of  a  base  and  is  not  in  directly  ionizable  condition  at  all. 


POTASSIUM  AND  AMMONIUM  679 

There  are  other  indications  that  the  amount  of  un combined  ammonia 
is  considerable.  Thus  the  organic  derivative  tetramethylammonium 
hydroxide  N(CH3)4OH  is  a  very  active  base  indeed,  and  one  of  the 
most  conspicuous  differences  between  it  and  ammonium  hydroxide 
is  that  it  cannot  decompose  into  water  and  a  non-ionizable  substance. 
It  is  all  available  for  ionization,  while  the  material  in  ammonia-water 
is  not.  It  is  also  noteworthy  that  ammonium  chloride  solution  is 
practically  neutral,  and  not  acid,  as  the  solution  of  the  chloride  of  a 
very  weak  base  would  be. 

Ammonium  Nitrate  NH^NO^.  —  This  is  a  white  crystalline 
salt  which  may  be  made  by  the  interaction  of  ammonium  hydroxide 
and  nitric  acid  (cf.  p.  395).  When  heated  gently  it  decomposes, 
giving  nitrous  oxide  and  water  (p.  539).  It  is  used  as  an  ingredient 
in  fireworks  and  explosives.  It  exists  in  no  fewer  than  four  solid 
physical  states.  The  melted  salt  solidifies  at  about  166°,  giving  crys- 
tals of  the  regular  system.  When  these  are  allowed  to  cool  somewhat, 
and  are  held  at  a  temperature  a  little  below  125.5°,  they  change 
gradually  into  a  mass  of  rhombohedral  crystals,  the  density  and  all 
other  physical  properties  altering  at  the  same  time.  This  tempera- 
ture is  a  transition  point  like  that  at  which  monoclinic  sulphur  assumes 
the  rhombic  form  (p.  412).  When  these  rhombohedral  crystals,  in 
turn,  are  held  at  a  temperature  a  little  below  83°  they  change  their 
form  once  more  into  crystals  which  belong  to  the  rhombic  system  and 
possess  a  third  distinct  set  of  physical  properties.  Finally,  below  35° 
a  fourth  change,  into  rhombic  needles,  takes  place,  and  this  condition 
of  the  substance  is  the  one  familiar  at  ordinary  temperatures.  All 
these  changes  proceed  in  the  reverse  order  when  the  temperature  is 
elevated  once  more. 

Ammonium  Carbonate.  —  When  ammonium  hydroxide  is 
treated  with  excess  of  carbon  dioxide  the  solution  gives,  on  evapora- 
tion, ammonium  bicarbonate  NH4HCO3.  This  is  a  white  crystalline 
salt  which  is  fairly  stable  at  the  ordinary  temperature.  It  has, 
however,  a  faint  odor  of  ammonia,  and  its  dissociation  becomes  very 
rapid  when  slight  heat  is  applied.  When  a  solution  of  this  salt  is 
treated  with  ammonium  hydroxide,  the  neutral  carbonate  is  formed: 

NH4HC03  +  NH4OH  <±  (NH4)2C03  +  H20. 

But  this  salt,  when  left  in  an  open  vessel,  loses  ammonia  very  rapidly, 
and  leaves  the  bicarbonate  behind. 


680  INORGANIC   CHEMISTRY 

The  substance  commonly  sold  as  ammonium  carbonate  is  the  so- 
called  sesquicarbonate,  and  is  made  by  sublimation  from  a  mixture  of 
ammonium  chloride  (or  ammonium  sulphate)  and  chalk  or  powdered 
limestone.  It  is  a  mixture,  in  approximately  equi-molar  proportions, 
of  ammonium  bicarbonate  and  ammonium  carbamate.  The  latter  is 
a  substance  related  to  urea,  and  formed  when  ammonia  and  carbon 
dioxide  gases  are  mixed  : 

NH2 


Ammonium  cyanate  is  interesting  on  account  of  its  rapid  trans- 
formation, when  warmed,  into  urea  (p.  583).  Ammonium  thiocyanate 
NH4NCS  is  a  white  salt  which  finds  some  application  in  analysis. 

Ammonium  Sulphate.  —  This  is  a  white  salt,  crystallizing  in 
rhombic  prisms,  which  is  used  chiefly  as  a  fertilizer.  By  electrolysis 
of  a  concentrated  solution  of  the  bisulphate,  ammonium  persulphate, 
which  is  less  soluble,  is  formed  and  crystallizes  out  (cf.  p.  449). 

Sulphides  of  Ammonium.  —  When  gaseous  hydrogen  sulphide 
and  ammonia  are  mixed  in  equi-molar  proportions  and  compressed  or 
strongly  cooled,  ammonium-hydrogen  sulphide  NH4HS  is  formed  as 
a  crystalline  deposit  on  the  vessel.  In  an  open  vessel,  at  the  ordinary 
temperature,  this  solid  dissociates  slowly  into  its  constituents.  The 
sulphide,  (NH4)2S,  can  be  produced  under  similar  conditions  by  using 
twice  as  much  ammonia.  But  it  is  much  less  stable  and  gives  up  half 
its  ammonia,  producing  the  acid  sulphide  very  quickly.  Solutions  of 
these  sulphides,  made  by  passing  hydrogen  sulphide  gas  into  ammo- 
nium hydroxide,  are  much  used  in  analysis.  The  sulphide  is  almost 
completely  hydrolyzed  by  water  into  the  acid  sulphide  and  ammonium 
hydroxide,  its  behavior  being  like  that  of  sodium  sulphide  (p.  418): 

2NH3  (dslvd)  +  H2S  (dslvd)  <=±  (NH4)2S  ^  2NH4+  +  S=)  <_  ™_ 

H20i^    OH-  +  H+J 


It  is  used  for  the  precipitation  of  sulphides,  such  as  zinc  sulphide, 
which  are  insoluble  in  water.  Although  the  S=  ions  are  not  numerous 
at  any  moment,  disturbance  of  the  equilibrium  by  their  removal,  when 
they  pass  into  combination,  causes  displacements  which  result  in  the 
generation  of  a  continuous  supply.  The  liquid  smells  strongly  of 
ammonia  and  hydrogen  sulphide  on  account  of  the  dissociation  of  the 
parent  molecules.  Because  of  this  dissociation,  the  salt  is  preferred 


POTASSIUM  AND  AMMONIUM  681 

to  potassium  or  sodium  sulphide  in  analysis.  The  excess  of  the 
reagent  can  be  driven  out  by  simply  boiling  the  mixture  for  a  few 
minutes,  all  of  the  above  equilibria  being  reversed.  Another  applica- 
tion in  analysis  depends  on  the  tendency  of  this  salt  to  unite  with 
certain  insoluble  sulphides,  particularly  those  of  tin,  arsenic,  and 
antimony  (q.v.),  giving  soluble  complex  salts. 

The  solution  dissolves  free  sulphur,  giving  yellow  polysulphides 
similar  to  those  of  potassium  (p.  674).  The  same  yellow  substances 
are  also  obtained  by  gradual  oxidation  of  ammonium  sulphide  when 
the  solution  of  this  salt  is  allowed  to  stand  in  a  bottle  from  which  the 
air  is  imperfectly  excluded. 

Microcosmic  Salt.  —  This  salt  would  be  named,  systematically, 
the  tetrahydrate  of  secondary  sodium-ammonium  orthophosphate 
NaNH4HP04,4H2O.  When  ammonium  chloride  and  ordinary  so- 
dium phosphate  are  mixed  in  strong  solution  the  hydrate  crystal- 
lizes out.  The  substance  is  used  in  bead  tests  (cf.  pp.  559,  560). 

Ammonium  Amalgam.  —  When  a  salt  of  ammonium  is  de- 
composed by  electrolysis  the  NH4  ion,  upon  its  discharge,  gives 
ammonia  and  hydrogen,  and  no  substance  NH4  is  obtained.  If,  how- 
ever, a  pool  of  mercury  is  used  as  the  negative  electrode,  the  NH4 
forms  an  amalgam  with  it,  and  there  seems  to  be  no  doubt  that  this 
substance  is  actually  present  in  solution  in  the  mercury.  While  the 
amalgam  is  being  formed  it  swells  up  and  gives  off  the  decomposition 
products  above  mentioned,  so  that  the  existence  of  the  substance  is 
only  temporary.  The  same .  material  may  be  obtained  by  putting 
sodium  amalgam  into  a  strong  solution  of  a  salt  of  ammonium.  The 
action  is  a  displacement  of  one  ion  by  another  (p.  403) : 

Na  (dslvd  in  mercury)  +  NH4+  — •>  NH4  (dslvd  in  mercury)  -f  Na+. 

This  behavior  is  interesting  since  it  is  in  harmony  with  the  idea  that 
ammonium,  if  it  could  be  isolated,  would  have  the  properties  of  a 
metal.  Substances,  other  than  metals,  are  not  miscible  with  mercury. 

Ammonium-ion  NH±+:  Analytical  Reactions.  —  Ionic  ammo- 
nium is  a  colorless  substance.  It  unites  with  negative  ions,  giving 
salts,  which,  in  the  majority  of  cases,  are  soluble.  Ammonium  chlo- 
roplatinate  (NH4)2PtCl6,  and  to  a  less  extent  ammonium  hydrogen 
tartrate  NH4HC4H4Oe,  are  insoluble  compounds,  and  their  precipita- 
tion is  used  as  a  test.  The  surest  means  of  recognizing  ammonium 


682  INORGANIC  CHEMISTRY 

compounds,  however,  consists  in  adding  a  soluble  base  to  the  sub- 
stance (cf.  p.  521).  The  ammonium  hydroxide,  which  is  thus  formed, 
gives  off  ammonia,  and  the  latter  may  be  detected  by  its  odor.  The 
quantity  of  the  ammonium  salt  present  may  be  determined  by  dis- 
tilling the  mixture  and  catching  the  distillate  in  a  measured  volume  of 
normal  hydrochloric  acid.  Determination  of  the  amount  of  the  acid 
remaining  unneutralized,  by  titration  with  a  standard  alkali  solu- 
tion, then  gives,  by  difference,  the  quantity  of  ammonium  hydroxide. 

Exercises.  —  1.   What  kind  of  metals  will,  in  general,  interact 
with  solutions  of  bases  (cf.  p.  646)? 

2.  Why  should  a  mixture  of  potassium  chlorate  and  antimony  tri- 
sulphide  be  explosive? 

3.  How  does  the  direct  vision  spectroscope  differ  from  the  ar- 
rangement here  described  (cf.  any  work  on  physics)? 

4.  Why  is  not  ammonium  carbamate  (p.  680)  formed  by  the 
neutralization  method? 

5.  How  should  you  set  about  making,  (a)  a  borate  of  potassium, 
(b)  potassium  pyrophosphate,  (c)  ammonium  nitrite,  (d)  ammonium 
chlorate,  (e)  ammonium  iodide? 

6.  Why  is  the  cleaning  of  platinum  wires,  as  usually  effected  by 
holding  them  in  the  Bunsen  flame,  assisted  by  periodical  dipping  into 
hydrochloric  acid  (p.  677)? 


CHAPTER  XXXIV 

SODIUM  AND   LITHIUM.     IONIC  EQUILIBRIUM   CONSIDERED 
QUANTITATIVELY 

SODIUM  chloride  forms  more  than  two-thirds  of  the  solid  matter 
dissolved  in  sea-water,  and  the  great  salt  deposits  are  largely  composed 
of  it.  Sea-plants  contain  sodium  salts  of  organic  acids,  just  as  land- 
plants  contain  potassium  salts.  Chile  saltpeter,  cryolite,  and  albite 
(a  soda  felspar)  are  important  minerals. 

Compounds  of  sodium  are  usually  cheaper  than  the  correspond- 
ing ones  of  potassium.  Also,  since  the  atomic  weight  of  sodium  is 
23,  against  39  for  potassium,  a  smaller  weight  of  the  sodium  com- 
pound will  produce  the  same  chemical  result.  For  these  two  reasons, 
sodium  compounds,  except  in  special  cases,  are  always  used  for  com- 
mercial purposes. 

Preparation.  —  Sodium  was  first  made  by  Davy  (1807)  by  elec- 
trolysis of  moist  sodium  hydroxide.  It  is  manufactured  by  the  elec- 
trolysis of  fused  sodium  hydroxide  by  a  method  in- 
vented by  Castner.  The  negative  electrode  projects 
through  the  bottom  of  the  iron  vessel  containing  the 
fused  hydroxide  (Fig.  142),  and  here  the  sodium  and 
hydrogen  are  liberated.  This  electrode  is  surrounded 
by  a  wire-gauze  partition  to  permit  circulation  in  the 
fused  mass,  but  prevent  escape  of  the  globules  of 
sodium.  This  is  surmounted  by  a  bell-shaped  vessel 
of  iron.  The  positive  electrode  is  an  iron  cylinder 
surrounding  the  gauze.  The  sodium  and  hydrogen  FlQ-  142- 
liberated  at  the  cathode,  being  lighter  than  the  fused  mass,  ascend 
into  the  iron  vessel,  under  the  edge  of  which  the  hydrogen  escapes. 
Oxygen  is  set  free  at  the  anode.  The  top  is  closed,  to  prevent  the 
sodium  from  burning.  The  melted  sodium  is  ladled  into  molds,  like 
candle  molds. 

Properties.  —  Sodium  is  a  soft,  shining  metal,  melting  at  97.5° 
and  boiling  at  742°.  The  green  vapor  is  a  monatomic  gas.  The 

683 


684  INORGANIC  CHEMISTRY 

metal  is  soluble  in  liquefied  ammonia,  giving  a  blue  solution.  The 
amalgam  with  mercury,  when  it  contains  more  than  a  small  amount 
of  sodium,  is  solid,  and  contains  one  or  more  compounds  of  the  two 
elements.  This  amalgam  is  often  used  instead  of  the  metal  sodium, 
since  the  dilution  or  combination  with  mercury  makes  the  inter- 
actions of  the  metal  more  easily  controllable.  Sodium  is  used  in  the 
manufacture  of  sodium  peroxide,  and  of  many  complex  carbon  com- 
pounds which  are  employed  as  drugs  and  dyes. 

Sodium  Hydride  NaH.  —  When  hydrogen  is  led  over  sodium 
at  364°,  in  such  a  way  that  the  upper  part  of  the  tube  is  cooler,  a 
matted  mass  of  fine  white  crystals  of  the  hydride  is  deposited  on  the 
cool  part  of  the  tube.  The  temperature  must  not  rise  beyond  430°, 
since  the  compound  dissociates  rapidly  at  this  temperature.  The 
properties  of  the  substance  are  similar  to  those  of  potassium  hydride 
(p.  663). 

Sodium  Chloride  NaCl.  —  Common  salt  is  obtained  from  the 
salt  deposits  of  Stassfurt,  Reichenhall  (near  Salzburg),  in  Cheshire, 
at  Syracuse  and  Warsaw  in  New  York,  at  Salina  in  Kansas,  in  Utah, 
California,  and  many  other  districts.  Natural  brines  are  obtained 
from  wells  in  various  parts  of  the  world.  Since  the  salt  can  seldom  be 
used  directly,  on  account  of  impurities  which  it  contains,  it  is  purified 
by  recrystallization  from  water.  Natural  brines,  which  are  some- 
times dilute,  are  often  concentrated  by  dripping  over  extensive  ricks 
composed  of  twigs.  When  the  resulting  brine  is  allowed  to  evaporate 
slowly  by  the  help  of  the  sun's  heat,  large  crystals,  sold  as  "  solar 
salt,"  are  obtained.  By  the  use  of  artificial  heat  and  stirring,  smaller 
crystals  of  greater  purity  can  be  secured.  In  northern  Russia,  the 
brine  is  allowed  to  freeze,  and  the  water  thus  removed  in  the  form  of 
ice  (p.  199).  Salt  intended  for  table  use  must  be  freed  from  the 
traces  of  magnesium  chloride  (q.v.)  present  in  the  original  brine  or 
deposit,  for  this  impurity  causes  it  to  absorb  moisture  more  vigor- 
ously from  the  air.  Addition  of  a  little  baking  soda  NaHCOs  remedies 
the  difficulty,  by  forming  the  insoluble  magnesium  carbonate.  The 
purest  salt  for  chemical  purposes  is  precipitated  from  a  saturated 
solution  of  salt  by  leading  into  it  hydrogen  chloride  gas.  Explana- 
tion of  this  effect  will  be  given  presently  (see  p.  699). 

Common  salt  crystallizes  in  cubes,  the  faces  of  which  are  usually 
hollow.  The  crystals  decrepitate  (p.  670)  when  heated,  and  melt  at 
about  820°.  Common  salt  is  the  source  of  all  sodium  compounds, 


SODIUM  AND  LITHIUM  685 

with  the  exception  of  the  nitrate.     From  it  come  also  most  of  the 
chlorine  and  hydrogen  chloride  used  in  commerce. 

The  Hydroxide  and  Oxides.  —  Sodium  hydroxide  NaOH  is 
prepared  both  by  the  action  of  slaked  lime  upon  sodium  carbonate 
and  by  the  electrolysis  of  a  solution  of  sodium  chloride,  precisely  as  is 
potassium  hydroxide  (p.  666). 

Sodium  hydroxide  is  a  highly  deliquescent  substance,  which,  when 
exposed  to  the  air,  first  liquefies  and  then  becomes  solid  on  account  of 
the  formation  of  sodium  carbonate.  Its  general  chemical  properties 
are  identical  with  those  of  potassium  hydroxide.  It  is  used  in  the 
manufacture  of  soap,  in  the  preparation  of  paper  pulp,  and  in  many 
other  chemical  industries. 

Sodium  peroxide  Na2O2  is  made  by  heating  sodium  at  300-400° 
hi  air  which  has  been  freed  from  carbon  dioxide.  The  sodium  is 
placed  on  trays  of  aluminium,  and  is  passed  into  the  furnace  against 
the  current  of  air.  In  this  way,  the  freshest  sodium  meets  the  air 
from  which  most  of  the  oxygen  has  been  removed,  and  the  action  is 
moderated.  Conversely,  the  almost  entirely  oxidized  sodium  meets 
the  freshest  air,  and  completion  of  the  oxidation  is  thus  assured. 
This  oxide  is  the  sodium  salt  of  hydrogen  peroxide.  When  thrown 
into  water,  it  decomposes  in  part,  in  consequence  of  the  heat  devel- 
oped, giving  sodium  hydroxide  and  oxygen.  With  careful  cooling, 
however,  much  of  it  can  be  dissolved.  By  interaction  with  acids 
it  yields  hydrogen  peroxide.  Sodium  peroxide  is  now  used  commer- 
cially for  oxidizing  and  bleaching  and,  in  the  form  of  oxone  (p.  85), 
as  a  source  of  oxygen.  The  ordinary  sodium  oxide  Na20  is  made  in 
the  same  way  as  is  potassium  oxide  (p.  668) . 

The  Nitrate  and  Nitrite.  —  The  occurrence  and  purification 
of  sodium  nitrate  NaNO3  have  already  been  described  (p.  525).  Its 
crystals  are  of  rhombohedral  form  (Fig.  69,  p.  173).  This  salt  is  one 
of  the  best  of  fertilizers,  since  it  furnishes  to  plants  the  nitrogen 
which  they  require  in  very  soluble  form.  It  is  used  also  in  the 
manufacture  of  potassium  nitrate,  of  nitric  acid,  and  of  sodium 
nitrite. 

Sodium  nitrite  NaN02  is  formed  by  heating  sodium  nitrate  with 
metallic  lead  and  recrystallizing  the  product  (p.  538).  Although  very 
soluble  it  is  less  so  than  potassium  nitrite,  and  is  therefore  more  easily 
prepared  in  pure  condition.  It  is  used  as  a  source  of  nitrous  acid  by 
manufacturers  of  organic  dyes. 


686 


INORGANIC   CHEMISTRY 


Manufacture  of  Sodium  Carbonate.  —  Natural  sodium  car- 
bonate is  found  in  Egypt  and  in  other  parts  of  the  world.  At  Owen's 
Lake,  California,  it  is  secured  by  solar  evaporation  of  the  water.  The 
sesquicarbonate  Na2CO3,NaHCO3,2H2O,  being  the  least  soluble  of 
the  carbonates  of  sodium,  is  the  one  deposited.  Locally,  small  quan- 
tities of  sodium  carbonate  are  still  made  by  the  burning  of  sea- weed. 
Up  to  the  close  of  the  eighteenth  century  this  was  the  only  source  of 
the  compound,  and  the  product  from  Spain,  known  commercially  as 
barilla,  was  ten  times  as  expensive  as  the  carbonate  now  is.  Hence 
glass  and  soap  were  proportionately  dearer  than  at  present. 

In  1791  the  French  Academy  offered  a  prize  for  the  discovery  of 
an  inexpensive  method  for  the  preparation  of  sodium  carbonate  from 
common  salt,  and  Le  Blanc  proposed  the  process  which  bears  his  name 


mis 


PlO.  143. 


and  is  still  in  use  in  two  factories  in  Europe.  During  the  Revolution 
his  factory  was  destroyed,  his  patents  were  declared  to  be  public 
property,  and  the  inventor  died  by  suicide.  The  chief  stages  of  Le 
Blanc's  process  involve  three  chemical  actions.  In  the  first  place, 
sodium  chloride  is  treated  with  an  equivalent  amount  of  sulphuric 
acid  in  a  large  cast-iron  or  earthenware  pan.  The  bisulphate  thus 
produced  (cf.  p.  206),  together  with  the  unchanged  sodium  chloride, 
is  raked  out  on  to  the  hearth  of  a  reverberatory  *  furnace  (Fig.  143) 
and  heated  more  strongly: 

NaCl  +  NaHS04  «=*  Na2S04  +  HC1  T . 

*  So  called  because  the  heated  gases  from  the  fire  are  deflected  by  the  roof  and 
play  upon  the  materials  spread  on  the  bed  of  the  furnace. 


SODIUM   AND  LITHIUM  687 

The  product  of  this  treatment  is  called  salt-cake.  The  second  and 
third  actions  which  follow  are  conducted  in  one  operation.  They  con- 
sist in  the  reduction  of  the  sodium  sulphate  by  means  of  powdered 
coal  and  the  interaction  of  the  resulting  sulphide  of  sodium  with  chalk 
or  powdered  limestone: 


4-  2C  -»  Na*S  +  2C02, 
Na2S  +  CaC03->  Na2C03  +  CaS. 

Formerly,  the  salt-cake,  limestone,  and  coal  were  stirred  upon  the 
hearth  of  a  reverberatory  furnace  and  worked  by  hand.  The  mate- 
rial was  collected  into  balls,  and  the  end  of  the  action  was  recognized 
by  the  fact  that  bubbles  of  carbon  monoxide  began  to  force  their  way 
to  the  surface  and  caused  little  jets  of  blue  flame.  The  gas  is  pro- 
duced by  the  action  of  the  coal  upon  the  calcium  carbonate,  excess  of 
both  of  these  substances  being  present  :  CaCO3  +  C  —  •>  CaO  -f  2CO. 
The  production  of  this  gas  gives  a  porous  texture  to  the  material, 
which  facilitates  the  solution  of  the  sodium  carbonate  in  the  final 
stage.  The  porous  product  is  called  black-ash.  In  modern  factories 
hand  labor  is  saved  by  giving  the  black-ash  furnace  the  form 
of  a  rotating  cylinder,  in  which  projections  from  the  walls  assist 
in  bringing  about  complete  mixing  of  the  materials  during  the 
action. 

The  black-ash  varies  very  much  in  composition.  It  commonly 
contains  45  per  cent  of  sodium  carbonate,  30  per  cent  of  calcium 
sulphide,  10  per  cent  of  calcium  oxide,  and  a  number  of  other  products 
and  impurities.  The  coal  used  in  the  operation  is  selected  so  as  to  be 
as  free  as  possible  from  combined  nitrogen,  the  presence  of  which 
leads  to  the  formation  of  cyanides. 

Calcium  sulphide  is  not  very  soluble  in  water,  and  is  but  slowly 
hydrolyzed  by  it  (p.  421),  especially  when  calcium  hydroxide  is 
present.  The  sodium  carbonate  is  therefore  extracted  from  the 
black-ash  by  water.  The  ash  is  placed  in  a  series  of  vessels  at  different 
levels,  and  a  stream  of  water  flows  from  one  vessel  to  another,  until, 
when  it  issues  from  the  last,  it  is  completely  saturated  with  sodium 
carbonate.  A  temperature  of  30°  to  40°,  at  which  the  solubility  of 
sodium  carbonate  is  at  a  maximum,  is  employed.  When  the  material 
in  the  first  of  the  vessels  has  been  exhausted,  the  water  is  allowed  to 
enter  the  second  vessel  directly,  and  a  vessel  containing  fresh  black- 
ash  is  added  at  the  lower  end  of  the  series.  In  this  way  the  most 
nearly  exhausted  ash  comes  in  contact  with  pure  water,  which  is  in 
the  best  position  to  dissolve  the  remaining  sodium  carbonate  rapidly, 


688  INORGANIC   CHEMISTRY 

while  the  fresh  black-ash  encounters  a  solution  already  almost  at  the 
point  of  saturation. 

The  saturated  solution  is  evaporated  in  shallow  pans  placed  in  the 
flues  of  the  furnaces,  and  the  monohydrate  Na2C03,H2O,  which  crys- 
tallizes from  the  hot  liquid,  is  raked  out  and  dried  by  heat,  leaving 
calcined  soda.  When  this  material  is  recrystallized  from  water  and  is 
allowed  to  deposit  itself  from  the  solution  at  the  ordinary  tempera- 
ture, the  decahydrate  Na2C03,10H20,  soda  crystals,  washing  soda, 
or  sal  soda  appears. 

The  Solvay,  or  ammonia-soda  process,  invented  in  1860,  has  now 
displaced  the  Le  Blanc  process.  It  differs  from  the  latter  by  in- 
volving almost  nothing  but  ionic  actions.  A  solution  of  salt  con- 
taining ammonia  and  warmed  to  40°  fills  a  tower  divided  by  a  number 
of  perforated  partitions.  Carbon  dioxide,  which  is  forced  in  below, 
makes  its  way  up  through  the  liquid.  The  ammonium  bicarbonate 
formed  by  its  action  undergoes  double  decomposition  with  the  salt, 
and  sodium  bicarbonate  which  is  precipitated  settles  upon  the  par- 
titions: 


NaCl  +  NH4HCO3  <=±NaHC04  +  NH4C1, 
or  HCO3~  +  Na+  <=»  NaHCO3  1  . 

The  solid  sodium  bicarbonate,  after  being  freed  from  the  liquid,  is 
heated  strongly  and  leaves  behind  sodium  carbonate: 

2NaHC03  ->  N^COs  +  H20t  +  C02t  . 

The  carbon  dioxide  which  is  liberated  passes  through  the  operation 
once  more.  The  supply  of  carbon  dioxide  is  generated  in  lime-kilns 
of  special  form.  The  mother-liquor  from  the  sodium  bicarbonate 
contains  ammonium  chloride.  This  is  decomposed  by  heating  with 
quicklime  from  the  kilns,  and  the  ammonia  which  is  thus  obtained  is 
available  for  the  treatment  of  another  batch.  A  solution  of  calcium 
chloride  is  thus  the  only  waste  product.  Some  of  this  is  used  for 
laying  dust  on  roads,  because,  being  a  deliquescent  salt,  the  moistened 
dust  does  not  dry  up. 

This  process  furnishes  a  much  purer  product  than  does  the  Le 
Blanc  process.  A  possible  rival  threatens  to  arise  in  the  treatment 
of  electrolytic  sodium  hydroxide  with  carbon  dioxide  gas. 

Properties  of  Sodium  Carbonate.  —  The  common  form  of 
sodium  carbonate  consists  of  large  monoclinic  crystals  of  the  decahy- 
tfrate  NaaCO^lQHgO,  This  substance  has  a  fairly  high  aqueous  ten- 


SODIUM  AND  LITHIUM  689 

sion,  and  loses  nine  of  the  ten  molecules  of  water  which  it  contains 
when  it  is  exposed  in  an  open  vessel.  When  warmed  it  melts  at  35.2°, 
giving  a  solution  of  sodium  carbonate  in  water.  The  residue  from 
evaporation,  above  35.2°,  is  the  monohydrate  Na2CO3,H2O.  At  higher 
temperatures,  or  with  low  atmospheric  aqueous  tension  (p.  152),  this 
in  turn  can  be  completely  dehydrated.  The  decahydrate  increases 
in  solubility  up  to  35.2°,  when  it  ceases  to  exist.  Just  above  this  tem- 
perature the  monohydrate  is  the  only  substance  which  is  stable.  The 
solubility  of  the  monohydrate  is  the  same  as  that  of  the  decahydrate 
at  35.2°,  and  diminishes  as  the  temperature  is  raised.  The  relations 
are  of  the  same  nature  as  in  the  case  of  sodium  sulphate  (p.  193).  In 
aqueous  solution,  sodium  carbonate  is  hydrolyzed  (3.2  per  cent  in 
Q.IN  sol.  at  25°),  and  shows  a  marked  alkaline  reaction.  The  com- 
pound is  used  in  large  amounts  for  the  manufacture  of  glass  and 
soap,  in  the  softening  of  water,  and  is  applied  in  innumerable  ways  in 
the  scientific  industries  for  purposes  akin  to  cleansing. 

All  the  familiar  salts  of  sodium,  excepting  sodium  nitrate  and 
the  peroxide,  are  made  by  the  treatment  of  sodium  carbonate  or 
sodium  hydroxide  with  acids. 

Sodium  Bicarbonate  NaHCOs  or  Baking  Soda.  —  This  salt 
is  formed  in  the  Solvay  process,  and  can  be  prepared  in  a  state  of 
purity  by  passing  carbon  dioxide  over  the  decahydrate  of  sodium 
carbonate : 

Na2C03,10H2O  +  C02<=»2NaHC03  +  9H2O. 

The  hydrate  is  spread  upon  a  grating,  through  which  the  water  gener- 
ated by  the  action  drips  away.  This  action  is  reversible,  and  sodium 
bicarbonate  shows,  even  in  the  cold,  an  appreciable  tension  of  carbon 
dioxide.  Even  a  solution  of  this  salt  gives  off  carbon  dioxide,  when 
boiled.  An  aqueous  solution  of  pure  sodium  bicarbonate  is  neutral 
to  phenolphthalem,  on  account  of  the  small  degree  of  ionization  of  the 
ion  HCO3~~.  Ordinarily,  however,  the  solution  is  alkaline,  on  account 
of  the  presence  of  the  carbonate,  which  is  hydrolyzed. 

Baking  Powders.  —  The  object  of  using  the  powder  is  to  gen- 
erate carbon  dioxide  in  the  dough.  The  bubbles  are  retained  because 
of  the  presence  of  the  sticky  gluten.  They  expand  when  the  dough 
is  heated  in  baking,  and  give  to  the  bread  its  porous  texture. 

Baking  soda,  alone,  will  give  off  carbon  dioxide,  but  the  sodium 
carbonate  which  it  leaves  behind  has  a  disagreeable  taste  and  acts 


690  INORGANIC  CHEMISTRY 

upon  the  gluten  causing  a  yellow  color  and  unpleasant  smell.  It 
also  tends  to  neutralize  the  acid  in  the  gastric  juice  and  so  impedes 
digestion.  To  prevent  this  result,  sour  milk  (containing  lactic 
acid)  and  even  vinegar  are  added.  Usually,  however,  a  baking 
powder  containing  an  acid  substance  along  with  the  bicarbonate  is 
employed.  Potassium  bitartrate  (cream  of  tartar)  KHC4H4O6  (p. 
608)  is  most  commonly  employed,  although  alum  and  primary  sodium 
orthophosphate  (p.  558)  are  also  used: 

HKC4H406  +  NaHC03  ->  NaKC4H406  +  H2C03  ->  H20  +  C02. 

The  cream  of  tartar  has  the  advantages  that  it  is  somewhat  in- 
soluble and  does  not  act  noticeably  upon  the  soda  before  the  mixing 
of  the  dough  is  complete,  and  that  the  sodium-potassium  tartrate 
(Rochelle  Salt)  produced  is  not  harmful.  A  little  starch  is  added 
to  baking  powders  to  keep  the  particles  of  the  two  other  ingredients 
apart,  and  prevent  gradual  interaction  before  use. 

For  raising  bakers'  bread,  yeast  is  employed,  and  time  is  allowed 
for  the  propagation  of  the  yeast  and  its  action  upon  the  sugar  (p. 
607)  in  the  flour.  A  little  molasses  or  malt  extract  is  often  added,  to 
ensure  a  sufficient  supply  of  sugar. 

The  whites  of  eggs  cause  cake  to  rise,  largely  because  they  are 
whipped  before  use,  and  bubbles  of  air,  which  expand  when  heated, 
are  thus  introduced. 

Sodium  Sulphate  7Va2SO4.  —  Anhydrous  sodium  sulphate 
(thenardite)  crystallizes  in  the  rhombic  system,  and  is  found  in  the 
salt  layers.  The  same  salt  is  contained  in  mineral  waters,  such  as 
those  of  Friedrichshall  and  Karlsbad.  It  is  formed  in  connection 
with  the  manufacture  of  nitric  acid  from  sodium  nitrate.  Some  of  it 
is  also  prepared  at  Stassfurt  by  dissolving  kieserite  MgSO4,H2O  in 
water  along  with  sodium  chloride.  When  the  solution  is  cooled  to  0°, 
the  decahydrate  of  sodium  sulphate  crystallizes  out,  and  magnesium 
chloride  remains  in  solution.  It  is  used,  as  a  substitute  for  sodium 
carbonate,  in  making  inexpensive  glass. 

The  decahydrate  of  sodium  sulphate  Na2S04, 10H20,  Glauber's 
salt,  forms  large  monoclinic  crystals  which  give  up  all  the  water  of  hy- 
dra tion  when  kept  in  an  open  vessel.  When  heated  the  crystals  melt 
at  32.4°,  and  are  resolved  into  the  sulphate  and  water.  The  relations 
of  the  hydrate  and  anhydrous  substance  in  respect  to  solubility  have 
been  fully  discussed  already  (p.  193).  When  the  decahydrate  is 
mixed  with  concentrated  hydrochloric  acid,  it  is  decomposed,  and  a 


SODIUM  AND  LITHIUM  691 

part  of  the  sulphate  is  converted  into  sodium  chloride,  the  second 
action  being  a  reversible  one.  This  is  one  of  those  actions  which 
proceed  spontaneously,  and  therefore  involve  a  diminution  in  the 
store  of  available  energy  in  the  system,  although,  so  far  as  heat  is 
concerned,  a  marked  absorption  of  this  form  of  energy  takes  place 
(cf.  p.  35).  The  combination  is  used,  in  fact,  as  a  freezing  mixture. 


Sodium  Thiosulphate  NazSzO^SHzO.  —  This  salt  is  made  by 
boiling  a  solution  of  sodium  sulphite  with  sulphur.  It  is  also  ob- 
tained by  boiling  sulphur  with  caustic  soda,  and  crystallizes  from  the 
mixed  solution: 

4S  +  GNaOH  ->  Na2S203  +  2Na2S  +  3H20. 

It  gives  large,  transparent  monoclinic  crystals  of  a  pentahydrate. 
When  heated  it  first  loses  the  water  of  hydration,  and  then  decom- 
poses, giving  sodium  sulphate,  which  is  the  most  stable  oxygen- 
sulphur  compound  of  sodium,  and  sodium  pentasulphide  : 

4Na2S2O3  ->  3Na2S04  +  Na2S5. 

From  the  latter,  four  unit-weights  of  sulphur  can  be  driven  by  stronger 
heating.  Sodium  thiosulphate  (hypo)  is  used  for  fixing  negatives 
in  photography  (q.v.),  and  by  bleachers  as  antichlor  (p.  448).  For 
sodium  hyposulphite  Na&C^,  see  p.  443. 

Phosphates  of  Sodium.  —  Common  sodium  phosphate  is  a  do- 
decahydrate  of  the  secondary  orthophosphate,  Na2HP04,12H2O.  It 
is  made  by  neutralization  of  phosphoric  acid  with  sodium  carbonate, 
and  crystallizes  from  the  solution  in  large,  transparent  monoclinic 
prisms.  Its  properties  have  already  been  discussed  (p.  559). 

Sodium  metaphosphate  NaPO3  is  used  for  bead  tests  (cf.  p.  560). 

Sodium  Tetraborate  Na2B^O79  or  Borax.  —  This  salt  com- 
bines with  ten  molecules  of  water,  forming  large,  transparent  prisms, 
Na2B4O7,10H2O.  When  heated  it  loses  water,  and  leaves  the  easily 
fusible  anhydrous  salt  in  glassy  form.  Its  sources  have  already  been 
discussed  under  borates  (p.  637).  It  is  used  as  an  ingredient  in  glazes 
for  porcelain,  in  soldering,  for  bead  reactions  (cf.  p.  640),  and  for 
preserving  foods. 

Sodium  Disilicate  No^Si^G*,.  —  This  salt  (p.  634)  is  used  for 
fireproofing  wood  and  other  materials  and  for  preserving  eggs.  Sand 


692  INORGANIC  CHEMISTRY 

which  is  moistened  with  it  and  pressed  in  molds,  forms,  after  baking, 
a  serviceable  artificial  stone.  Since  silicic  acid  is  a  feeble  acid,  this 
salt  is  hydrolyzed,  and  gives  a  strongly  alkaline  solution  (p.  648). 
For  sodium  cyanide,  see  p.  721. 

Properties  of   Sodium-ion  Na+:    Analytical  Reactions.  — 

Sodium-ion  is  a  colorless  ionic  material  which  unites  with  all  negative 
ions.  Practically  all  the  salts  so  formed  are  soluble  in  water.  The 
only  ones  which  can  be  precipitated  are  sodium  fluosilicate  Na2SiF6, 
made  by  the  addition  of  hydrofluosilicic  acid  to  a  strong  solution  of 
a  sodium  salt,  and  sodium-hydrogen  pyroantimoniate  Na2H2Sb2O7, 
made  by  similar  addition  of  the  corresponding  potassium  salt.  The 
two  yellow  lines  in  the  spectrum  are  characteristic.  If  the  yellow 
light  persists  longer  than  could  be  accounted  for  by  the  ordinary 
deposit  of  dust  on  the  wire,  a  sodium  compound  is  present  in  the 
material. 

LITHIUM 

The  compounds  of  lithium  are  made  from  amblygonite,  a  mixed 
phosphate  and  fluoride  of  aluminium  and  lithium.  It  occurs  in 
lepidolite  (a  lithia  mica)  and  in  other  rare  minerals.  Traces  of  com- 
pounds of  the  element  are  found  widely  diffused  in  the  soil,  and  are 
taken  up  by  plants,  particularly  tobacco  and  beets,  in  the  ashes  of 
which  the  element  may  be  detected  spectroscopically. 

The  metal  is  liberated  by  electrolysis  of  the  fused  chloride,  the 
manipulation  being  facilitated  by  the  addition  of  some  potassium  chlo- 
ride to  lower  the  melting-point  of  the  lithium  salt.  The  melting- 
point  and  boiling-point  of  the  free  element  are  higher  than  those  of 
any  other  alkali-metal,  and  the  specific  gravity  (0.53)  is  lower  than 
that  of  any  other  metal  whatever.  Lithium  not  only  floats  upon 
water,  but  also  in  the  petroleum  in  which  it  is  preserved. 

The  metal  behaves  towards  water  and  oxygen  like  sodium  (p.  115). 
It  unites  directly  and  vigorously  with  hydrogen  (LiH),  nitrogen 
(Li3N),  and  oxygen  (Li2O),  forming  stable  compounds.  The  chloride 
crystallizes  in  octahedra  (p.  172).  The  relative  insolubility  (Table) 
of  the  hydroxide  LiOH,  the  carbonate  Li2COs,  and  the  phosphate 
Li3PO4,2H2O,  is  in  sharp  contrast  to  the  easy  solubility  of  the  cor- 
responding compounds  of  the  other  alkali-metals,  and  links  lithium 
with  magnesium.  The  compounds  of  lithium  give  a  bright-red  color  to 
the  Bunsen  flame,  and  a  bright-red  and  a  somewhat  less  bright  orange 
line  are  seen  in  the  spectrum.  The  carbonate  is  used  in  medicine. 


IONIC  EQUILIBRIUM  693 

IONIC  EQUILIBRIUM,  CONSIDERED  QUANTITATIVELY 

The  Simple  Case.  —  In  view  of  the  predominance  of  ionic 
actions  in  the  chemistry  of  the  metals,  and  of  the  determinative  effect 
of  ionic  equilibria  on  many  actions,  it  is  essential  that  we  should  be 
prepared  in  future  for  a  more  exact  study  of  these  phenomena  than 
we  have  hitherto  attempted.  The  whole  basis  for  this  exact  study 
has  already  been  supplied,  and  only  more  specific  application  of  the 
principles  is  demanded.  The  basis  referred  to,  which  should  now 
be  re-read  as  a  preliminary  to  what  follows,  is  contained  in,  (1)  the 
discussion  of  chemical  equilibrium  in  general  (pp.  287-295),  (2)  the 
application  of  the  same  principles  to  ionic  equilibrium  (p.  359), 
and  (3)  the  illustration  of  this  application  in  the  case  of  cupric 
bromide  (pp.  378-380). 

In  the  first  place,  the  principles  themselves  must  be  recalled. 
When  acetic  acid,  for  example,  is  dissolved  in  water,  it  is  ionized  thus : 

HC2H3O2  *±  H+  +  C2H302-. 

The  amount  of  molecular  acetic  acid  dissociated  per  second  in  a  given 
amount  of  the  solution  is  proportional  to  the  concentration  of  the 
molecules,  while  the  amount  of  the  two  ionic  materials,  hydrogen-ion 
and  acetate-ion,  uniting  to  form  molecules  of  acetic  acid  depends  on 
the  frequency  of  the  encounters  of  the  two  kinds  of  ions  and  is  propor- 
tional to  the  ionic  concentrations  (p.  358).  The  unit  of  concentration 
(p.  294)  is  1  mole  per  liter,  or,  in  the  present  case,  60  g.  of  the  acid, 
1  g.  of  hydrogen-ion,  and  59  g.  of  acetate-ion,  respectively,  per  liter, 
for  these  numbers  represent  the  weight  of  one  mole  of  each  compo- 
nent. According  to  the  law  of  concentration  (p.  297) : 

[H+]  X  [C2H302-]  _  v 
[HC2H302] 

and  the  numerical  value  of  this  fraction,  or  of  K,  remains  unchanged 
whatever  the  total  concentration  of  the  solution  may  be.  If  the  solu- 
tion is  diluted,  for  example,  [H+]  and  [C2H3O2~]  diminish  relatively 
less  quickly  than  [HC2H3O2]  in  order  that  the  value  of  the  whole 
expression  may  remain  the  same.  This  is  accomplished  by  ioniza- 
tion  of  a  part  of  the  material  whose  concentration  is  [HC2H3O2]  and 
its  transference  to  the  ionic  forms  whose  concentrations  are  [H+]  and 
[CjHsOr],  respectively  (p.  299). 

A  numerical  example  will  show  that  this  law  of  concentration  ex- 
presses the  facts  with  considerable  exactness.  The  data  in  regard  to 
acetic  acid  are  as  follows  (p.  365) : 


694 


INORGANIC  CHEMISTRY 


Acetic  Acid. 

I. 

Molar  Con- 
centration. 

II. 

Proportion 
of  Whole  Mate- 
rial Ionized. 

III. 

[H+J  and  IC2H3O2~] 
(I  X  II). 

Uni-molar     .    .    . 

1.0 

0.004 

0.004 

1.0  -0.004 

Deci-molar   .   .   . 

0.1 

0.013 

0.0013 

0.1  -0.0013 

Centi-molar  .    .    . 

0.01 

0.0407 

0.000407 

0.01-0.000407 

- 

Now  [H+]  =  [C2H3O2  ],  since  the  ions  are  produced  in  equal  numbers. 
Also,  for  our  purpose,  the  numbers  to  be  subtracted  in  column  iv  are 
relatively  so  small  that  the  values  1,  0.1,  and  0.01  may  be  taken  to 
represent  [H^flsC^]  without  appreciable  error.  Hence,  substituting 
the  data  in  equation  (1)  above,  we  have: 


(.004)' 


=  .0460. 


(.0013)' 
.1 


=  .0469. 


(.000407)- 
.01 


=  .04165. 


In  other  words,  although  the  last  solution  is  a  hundred  times  more 
dilute  than  the  first,  and  the  degree  of  ionization  has  increased  ten 
times,  the  whole  expression  remains  close  to  the  value  0.04165  and  is 
essentially  constant. 

When  conductivity  data,  like  the  above,  are  applied  in  the  same 
way  to  the  cases  of  more  highly  ionized  substances,  the  values  of  K 
are  less  nearly  constant.  It  is  supposed  that  with  this  class  of  sub- 
stances the  measurements  of  degrees  of  ionization  by  the  conductivity 
method  are  less  accurate,  although  the  cause  of  the  discrepancy  has 
not  been  fully  determined.  However,  in  the  very  general  applica- 
tions of  the  data,  which  are  all  that  we  shall  be  required  to  make,  the 
conclusions  will  not  be  affected  by  this  fact. 

Excess  of  One  Ion.  —  In  the  case  of  cupric  bromide  (p.  378), 
we  showed  that  increasing  the  concentration  of  the  bromide-ion  dis- 
placed the  equilibrium  by  favoring  the  union  of  the  ions  to  form  molec- 
ular cupric  bromide :  2Br~  +  Cu++  — >  CuBr2.  This  we  speak  of  as 
a  repression  of  the  ionization  of  the  cupric  bromide.  Now,  if  the 
substance  is  a  slightly  ionized  one,  like  a  weak  acid  or  a  weak  base,  the 
repression  of  the  ionization  through  the  formation  of  molecules  in 
this  way  may  remove  so  many  of  that  one  of  the  ions  which  is  not 
present  in  excess  (corresponding  to  the  Cu++  in  the  foregoing  illus- 
tration), that  the  mixture  will  no  longer  respond  to  tests  for  the  ion 
so  removed.  This  is  an  interesting  and  very  common  case.  The 


IONIC  EQUILIBRIUM  695 

behavior  of  acetic  acid,  a  weak,  slightly  ionized  acid,  will  serve  as  an 
illustration. 

In  normal  solution  (60  g.  in  1  1.)  acetic  acid  is  only  .004  ionized 
(p.  694),  so  that,  in  the  equation  for  the  equilibrium, 

(.996)  HC2H302  <=»  H+  (.004)  +  C2H302-  (.004), 

the  relative  proportions  are  as  shown  by  the  numbers  in  parenthesis. 
If  the  whole  of  the  acid  (60  g.)  were  ionized,  there  would  be  1  g.  of 
hydrogen-ion  per  liter.  Yet,  even  in  this  much  smaller  concentration 
(.004  g.  per  liter),  the  acid  taste  of  the  H+  and  its  effect  upon  indi- 
cators can  be  distinctly  recognized.  If,  now,  solid  sodium  acetate  is 
dissolved  in  the  solution,  the  liquid  no  longer  gives  an  add  reaction 
with  one  of  the  less  delicate  indicators,  like  methyl  orange  (q.v.). 
The  explanation  is  simple.  Sodium  acetate  is  highly  ionized.  It 
gives,  therefore,  a  large  concentration  of  acetate-ion  to  a  liquid  for- 
merly containing  very  little.  This  causes  a  greatly  increased  union 
of  the  H+  ions  and  C2H3O2~  ions,  and  the  former,  being  already  very 
few  in  number,  disappear  almost  entirely.  Hence  the  solution  be- 
comes, to  all  intents  and  purposes,  neutral.  There  is  no  less  acetic 
acid  present  than  before,  but  the  concentration  of  hydrogen-ion  is 
very  much  smaller. 

Formulation  and  Quantitative  Treatment  of  the  Case  of 
Excess  of  One  Ion.  —  If  the  semi-mathematical  mode  of  formu- 
lating an  equilibrium  (p.  359),  as  applied  to  the  case  of  an  ionogen,  be 
employed  here,  the  foregoing  general  statements  may  be  made  more 
precise  and  the  conclusions  clearer.  The  value  of  K  is  constant, 
whether  the  strength  of  the  solution  of  acetic  acid  is  great  or  small, 
and  even  when  another  substance  with  a  common  ion  is  present.  In  the 
latter  case,  [C2H302~]  and  [H+]  stand  for  the  whole  concentrations  of 
each  of  these  ionic  substances  from  both  sources. 

Now,  as  we  have  seen,  in  normal  acetic  acid  [H+]  =  .004, 
[C2H3O2~]  =  .004  (for  the  number  of  each  kind  of  ions  is  the  same), 
and  [HC2H302]  =  .996,  practically  1.  Substituting  in  the  formula 

0.004  X  0.004 


When  sodium  acetate  is  dissolved  in  the  liquid  until  the  solution  is 
normal  in  respect  to  this  substance  also,  the  following  additional 
equilibrium  has  to  be  considered: 

(.47)  NaC2H302  ^±  Na+  (.53)  +  C^OZ~  (.53). 


696  INORGANIC   CHEMISTRY 

The  concentration  of  acetate-ion  from  this  source  is  .53,  so  that,  in  the 
mixture  of  acid  and  salt,  the  concentration  of  acetate-ion  [C2H302~] 
will  be  .53  +  .004  =  .534,  or  nearly  134  times  larger  than  in  the  acid 
alone.  Hence,  in  order  that  the  product  [H+]  X  [C2H302~]  may  re- 
cover, as  it  must,  a  value  much  nearer  to  the  old  one,  [H+]  must  be 
diminished  to  something  like  y^  of  its  former  magnitude.  That 

is,   [H+]   will  become  equal  to  about  0.00003,  0-00003  x  °-534  = 

K  (=0.0416),  the  rest  of  the  hydrogen-ion  uniting  with  a  correspond- 
ing amount  of  the  acetate-ion  to  form  molecular  acetic  acid.  The 
effect  of  adding  this  amount  of  sodium  acetate  therefore  is,  as  we  have 
seen,  to  reduce  the  concentration  of  the  hydrogen-ion  below  the 
amount  which  can  be  detected  by  use  of  an  indicator  like  methyl 
orange. 

This  effect  is  of  course  reciprocal,  and  the  ionization  of  the  sodium 
acetate  will  be  reduced  also.  But  the  acetate-ion  furnished  by  the 
acetic  acid  is  relatively  so  small  in  amount  (.00003  against  .53)  that 
the  effect  it  produces  on  the  ionization  of  the  salt  is  imperceptible. 

It  will  be  noted  that  the  acetate-ion  and  hydrogen-ion  disappear 
in  equivalent  quantities,  for  they  unite.  There  is,  however,  so  much 
of  the  former  that  the  loss  it  suffers  goes  unremarked,  while  there  is  so 
little  of  the  latter  that  almost  none  of  it  remains.  When  substances 
of  more  nearly  equal  degrees  of  ionization  are  used,  both  effects  are 
equally  inconspicuous.  Thus,  sodium  chloride  and  hydrogen  chloride 
in  normal  solutions  yield  approximately  equal  concentrations  of 
chloride-ion  (.784  and  .676).  Hence,  if  one  mole  of  sodium  chloride 
were  to  be  dissolved  in  the  portion  of  water  already  containing  one 
mole  of  hydrogen  chloride,  the  concentration  of  the  chloride-ion,  at  a 
very  rough  estimate,  would  be  nearly  doubled.  If  this  doubling  of  the 
concentration  of  chloridion  almost  halved  that  of  the  hydrogen-ion 
(.784),  in  order  that  the  expression  [Cl~]  X  [H+]  -r-  [HC1]  might  re- 
main constant,  the  concentration  of  the  hydrogen-ion  would  still  be 
about  .400  and  therefore  100  times  as  great  as  in  molar  acetic  acid. 
It  is  thus  altogether  impossible  to  reduce  the  concentration  of  the 
hydrogen-ion  given  by  an  active  acid  like  hydrochloric  acid  below  the 
limit  at  which  indicators  are  affected,  for  there  is  no  way  of  introduc- 
ing the  enormous  concentration  of  the  other  ion  which  the  theory 
demands. 

With  more  crude  means  of  observation  than  indicators  afford, 
effects  like  this  last  may  sometimes  be  rendered  visible.  This  was 
the  case  with  cupric  bromide  solution,  to  which  potassium  bromide 


IONIC  EQUILIBRIUM  697 

was  added  (p.  379).  The  blue  of  the  cupric-ion  disappeared  from 
view,  while  much  cupric-ion  was  still  present,  because  the  brown 
color  of  the  molecular  cupric  bromide  covered  it  up  completely. 

Special  Case  of  Saturated  Solutions.  —  The  commonest  as 
well  as  the  most  interesting  application  of  the  conceptions  developed 
above  is  met  with  in  connection  with  saturated  solutions,  especially 
those  of  relatively  insoluble  substances. 

The  situation  in  a  system  consisting  of  the  saturated  solution  and 
excess  of  the  solute  has  been  discussed  already  (read  p.  382).  In  the 
case  of  potassium  chlorate,  for  example,  we  have  the  following  scheme 
of  equilibria : 

KC103  (solid)  <=»  KC103  (dslvd)  <=>  K+  +  ClOf . 

Solution  of  the  solid  is  promoted  by  the  solution  pressure  of  the  mole- 
cules, while  it  is  opposed  by  the  diffusion  pressure  of  the  dissolved 
substance,  and  the  solution  is  saturated  when  these  tendencies  produce 
equal  effects  (p.  186).  Now  it  must  be  noted  that  the  tendency  di- 
rectly opposed  to  the  solution  pressure  is  the  partial  diffusion  pressure 
of  the  dissolved  molecules  alone.  The  chief  contents  of  the  solution, 
the  molecules  and  two  kinds  of  ions  of  the  salt,  and  any  foreign  mate- 
rial that  may  be  present,  are  like  a  mixture  of  gases,  and  the  principle 
of  partial  pressure  (p.  Ill)  is  to  be  applied.  The 
ions  and  the  foreign  material  do  not  deposit  them- 
selves upon  the  solid,  and  take,  therefore,  no  part 
directly  in  the  equilibrium  which  controls  solubility. 
In  respect  to  this  the  ions  are  themselves  foreign 
substances.  Hence  the  conclusion  may  be  stated 
that,  in  solutions  saturated  at  a  given  temperature 
by  a  given  solute,  the  concentration  of  the  dissolved  n  KCIOi 

molecules  of  the  solute  considered  by  themselves 
will  be  constant  whatever  other  substances  may  be 
present. 

The  total  "solubility"  of  a  substance,  as  we 
have  used  the  term  hitherto,  is  made  up  of  a  mo-  Fia*  144< 

lecular  and  an  ionic  part.  The  latter,  as  we  shall  presently  see,  is 
not  constant  when  a  foreign  substance  containing  a  common  ion  is 
already  in  the  liquid.  Since  the  treatment  of  the  subject  requires  us 
now  to  distinguish  between  the  two  portions  of  the  solute,  a  diagram 
(Fig.  144)  will  assist  in  emphasizing  the  distinction.  The  material  at 
the  bottom  is  the  salt.  The  molecules  and  ions  are  to  be  thought 


mK++mCIO.~ 


698  INORGANIC  CHEMISTRY 

of  as  being  mixed  and  as  being  present  in  numbers  represented  by 
the  factors  n  and  ra.  Since  no  foreign  body  is  present,  the  two  ions 
in  this  case  are  equal  in  number. 

When  we  now  apply  these  ideas  to  the  mathematical  expression  of 
the  relation: 

.[K+]X  [CIO,"] 
[KC1O,] 

we  perceive  that,  in  a  saturated  solution,  [KClOs],  the  concentration 
of  the  molecules,  is  constant.  Transposing,  we  have 

[K+]  X  [CIO,-]  =  K  [KC10,]  =  K'. 

Hence  the  relation  leads  to  the  important  conclusion  that,  in  a  satu- 
rated solution  the  product  of  the  molar  concentrations  of  the  ions 
is  constant.*  This  product  is  called  the  ion-product  constant  for  the 
substance.  The  law  of  the  constancy  of  the  ion-product  in  a  satu- 
rated solution  is  one  of  the  most  useful  of  the  principles  of  chemistry. 
It  enables  us  to  explain  all  the  varied  phenomena  of  precipitation  and 
of  the  solution  of  precipitates  in  a  consistent  manner.  These  appli- 
cations of  the  principle  will  be  explained  in  the  next  chapter.  One 
curious  kind  of  precipitation  will  be  described  here,  however,  as  an 
illustration  of  the  use  of  the  principle. 

Illustration  of  the  Principle  of  Ion-Product  Constancy.  — 

When,  to  a  saturated  solution  of  one  of  the  less  soluble  salts,  a  strong 
solution  of  a  salt  having  one  ion  in  common  with  the  first  salt  is  added, 
precipitation  of  the  first  salt  frequently  takes  place.  This  happens, 
for  example,  with  a  saturated  solution  of  potassium  chlorate,  which 
is  not  very  soluble  (molar  solubility  0.52,  see  Table).  The  concentra- 
tions [K+]  and  [ClOs~]  being  small,  one  may  easily  increase  the  value 
for  one  of  the  ions,  say  [C1O3~],  fivefold,  by  adding  a  chlorate  which 
is  sufficiently  soluble.  To  preserve  the  value  of  the  product  [K+]  X 
[C103~],  the  value  of  [K+]  will  then  have  to  be  diminished  at  once  to 
one-fifth  of  its  former  value.  This  can  occur  only  by  union  of  the 
ionic  material  it  represents  with  an  equivalent  amount  of  that  for 

*  The  principle  of  constant  concentration  of  dissolved  molecules,  stated  above, 
has  been  shown  to  express  the  facts  very  inaccurately.  Now  the  principle  of  the 
constancy  of  the  ratio  of  the  ion-product  to  the  concentration  of  the  molecules  is 
also  inaccurate  in  the  case  of  highly  ionized  substances,  yet  in  such  a  way  that  the 
two  errors  neutralize  one  another.  Thus,  the  principle  of  ion-product  constancy 
here  given  is  in  itself  fairly  exact. 


IONIC  EQUILIBRIUM  699 

which  [C103~  ]  stands.  The  molecular  material  so  produced  will  thus 
tend  at  first  to  swell  the  value  of  [KC103].  But  the  value  of  [KC103] 
cannot  be  increased,  for  the  solution  is  already  saturated  with  molecules, 
so  that  the  new  supply  of  molecules,  or  others  in  equal  numbers, 
will  be  precipitated.  Hence  the  ionic  part  of  the  dissolved  substance 
may  be  diminished,  the  equilibria  (p.  697)  may  be  partially  reversed, 
and  we  may  actually  precipitate  a  part  of  the  olissolved  material 
without  introducing  any  substance,  which,  in  the  ordinary  sense,  can 
interact  with  it. 

In  point  of  fact,  when,  to  a  saturated  solution  of  potassium  chlo- 
rate there  is  added  a  saturated  solution  of  potassium  chloride  KC1  or 
of  sodium  chlorate  NaClO3,  a  precipitate  of  potassium  chlorate  is 
thrown  down.  These  two  salts,  containing  each  one  of  the  ions  of 
KC1O3,  and  being  much  more  soluble  than  the  latter  (see  Table), 
increase  the  concentration  of  one  ion  and  cause  the  precipitation  in 
the  fashion  just  explained. 

The  product  of  the  concentrations  of  the  ions,  for  example  [K+ ]  X 
[C1O3~],  is  called  also  the  solubility  product,  because  these  two  values 
jointly  determine  the  magnitude  of  the  solubility  of  the  substance. 
The  solubility  of  the  molecules  is  irreducible,  but  the  ionic  part  of  the 
dissolved  material  may  become  vanishingly  small  if  the  value  of  either 
[X+]  or  [Y~]  is  very  minute.  The  ionic  part  of  any  particular  sub- 
stance is  made  up  of  the  smaller  of  the  two  concentrations  of  the  ionic 
substances  which  it  yields,  plus  an  equivalent  amount,  and  no  more, 
of  the  concentration  of  the  other  ion.  The  rest  of  the  other  ionic 
substance  is  part  of  the  solubility  of  some  other  component. 

Other  Illustrations.  —  The  precipitation  of  sodium  chloride 
from  a  saturated  solution,  by  the  introduction  of  gaseous  hydrogen 
chloride  (p.  684),  is  to  be  explained  in  the  same  way.  The  equilibria: 

NaCl  (solid)  ?±  NaCl  (dslvd)  <=»  Na+  +  Cl- 
are reversed  by  the  introduction  of  additional  Cl~~  from  the  very 
soluble,  and  highly  ionized  HC1. 

A  steady  stream  of  hydrogen  chloride  is  often  obtained  by  drop- 
ping concentrated  sulphuric  acid  into  saturated  hydrochloric  acid: 

H+  +  Cr  ?±  HC1  (dslvd)  <=±  HC1  (gas). 

The  effect  is  due  in  part  to  repression  of  the  ionization  of  the  hydro- 
gen chloride  and  elimination  of  molecules  of  the  gas  from  the  water, 
which  is  already  saturated  with  molecules  of  the  same  kind. 


700  INORGANIC   CHEMISTRY 

The  formation  of  potassium  hydroxide  (p.  666)  ceases  when  a 
certain  concentration  has  been  reached.  This  occurs  because  the 
concentration  of  OH~,  which  rapidly  increases,  is  a  factor  in  the 
solubility  product  of  calcium  hydroxide,  [Ca++]  X  [OH~]2.  With 
much  OH~,  little  Ca++  is  required  to  give  the  constant,  numerical 
value  of  the  product.  When  the  concentration  [Ca++]  from  the 
hydroxide  has  become  about  as  small  as  that  from  the  carbonate, 
the  motive  for  the  interaction  has  been  removed.  This  principle 
is  thus  as  important  in  industrial  operations  as  it  is  in  analytical 
and  other  laboratory  experimentation. 

Exercises.  —  1.  The  vapor  density  of  sodium  peroxide  has  not 
been  determined.  Why  is  the  formula  NaaC^  assigned  to  it? 

2.  Construct  a  scheme  of  equilibria  (p.  687)  showing  the  hydroly- 
sis of  calcium  sulphide.     Why  does  the  presence  of  calcium  hydroxide 
diminish  the  tendency  to  hydrolysis? 

3.  What  will  be  the  effect  of  adding  a  concentrated  solution  of 
silver  nitrate  to  a  saturated  solution  of  silver  sulphate  or  of  silver 
acetate  (see  Table  of  solubilities)? 

4.  Although  a  20  per  cent  solution  of  soap  can  easily  be  made, 
a  0.5  per  cent  solution  can  be  salted  out  (p.  623).     How  does  this 
fact  show  that  salting  out  is  not  an  operation  like  the  precipitations 
just  discussed? 

5.  One  equivalent  or  less  of  concentrated  sulphuric  acid  interacts 
in  the  cold  with  sodium  carbonate  to  give  normal  sodium  sulphate, 
and  with  sodium  chloride  to  give  sodium-hydrogen  sulphate  (p.  206). 
Why  this  difference?     Which  of  the  two  products  would  be  obtained 
by  the  action  of  cold  concentrated  sulphuric  acid  on,  (a)  sodium 
sulphite,  (6)  sodium  nitrate,  (c)  sodium  nitrite? 


CHAPTER   XXXV 
METALLIC  ELEMENTS   OF   THE  ALKALINE  EARTHS 

CALCIUM,  STRONTIUM,  BARIUM 

The  Chemical  Relations  of  the  Elements.  —  The  metals  of 
this  group,  calcium  (Ca,  at.  wt.  40.07),  strontium  (Sr,  at.  wt.  87.63), 
and  barium  (Ba,  at.  wt.  137.37),  constitute  a  typical  chemical  family, 
both  in  the  qualitative  resemblance  to  one  another  of  the  elements  and 
of  the  corresponding  compounds,  and  in  the  quantitative  variation  in 
the  properties  with  increasing  atomic  weight.  The  metals  them- 
selves displace  hydrogen  vigorously  from  cold  water,  giving  hydrox- 
ides. The  solutions  of  these  hydroxides,  although  dilute,  on  ac- 
count of  a  rather  small  solubility,  are  strongly  alkaline  in  reaction. 
The  high  degree  of  ionization  of  the  hydroxides  recalls  the  hydrox- 
ides of  the  metals  of  the  alkalies,  and  their  relative  insolubility  the 
hydroxides  of  the  "earths"  (q.v.). 

In  all  their  compounds,  calcium,  strontium,  and  barium  are  biva- 
lent. The  hydroxides  are  formed  by  union  of  the  oxides  with  water 
and  are  progressively  less  easy  to  decompose  by  heating,  barium 
hydroxide  being  the  hardest.  The  carbonates,  when  heated,  yield  the 
oxide  of  the  metal  and  carbon  dioxide,  barium  carbonate  being  the 
most  difficult  to  decompose.  The  nitrates,  when  heated  moderately, 
give  the  nitrites,  but  the  latter  are  broken  up  by  heating  and  yield 
the  oxide  of  the  metal,  and  nitrogen  tetroxide.  In  these  and  other 
respects  the  compounds  of  the  metals  of  the  alkaline  earths  resemble 
those  of  the  heavy  metals  and  differ  from  those  of  the  metals  of  the 
alkalies.  Barium  approaches  the  latter  most  nearly. 

The  table  of  solubilities  shows  that  the  chlorides  and  nitrates  of 
calcium,  strontium,  and  barium  are  all  soluble  in  water,  the  solubility 
diminishing  in  the  order  given.  The  sulphates  and  hydroxides  cover 
a  wide  range  from  slight  solubility  to  extreme  insolubility.  Of  the 
sulphates,  2000,  110,  and  2.3  parts,  respectively,  dissolve  in  one 
million  parts  of  water.  In  the  case  of  the  hydroxides  the  order  of 
magnitudes  is  reversed,  and  the  corresponding  numbers  are  1700, 
7700,  and  37,000.  The  carbonates  are  almost  as  insoluble  as  is 

701 


702  INORGANIC  CHEMISTRY 

barium  sulphate.     The  element  radium  (Ra,  at.  wt.  226),  belongs  to 
this  family  (see  under  Uranium). 

CALCIUM  CA 

Occurrence.  —  The  fluoride,  and  the  various  forms  of  the  car- 
bonate, sulphate,  and  phosphate  which  are  found  in  nature,  are 
described  below.  As  silicate,  calcium  occurs,  along  with  other  metals, 
in  many  minerals  and  rocks.  It  is  found  also  in  plants,  and  its  com- 
pounds are  important  constituents  of  the  bones  and  shells  of  animals. 

The  Metal.  —  Although  the  alkali  metals  can  be  liberated  by 
heating  the  carbonates  with  carbon,  the  metals  of  the  present  family 
are  not  obtainable  by  this  means.  This  may  be  due,  in  part,  to 
imperfect  contact  between  the  materials  in  consequence  of  the  in- 
fusibility  of  the  oxides.  Calcium  is  most  easily  made  by  electrolysis 
of  the  molten  chloride,  to  which  calcium  fluoride  is  added  to 
lower  the  melting-point.  A  hollow  cylinder,  made  of  blpcks  of 
carbon  bolted  together  and  open  above,  forms  the  anode.  A  rod  of 
copper  hanging  so  that  its  end  dips  into  the  melt  forms  the  cathode. 
The  melting  of  the  anhydrous  calcium  chloride  with  which  the 
cylinder  is  filled  is  started  by  means  of  a  thin  rod  of  carbon  laid 
across  from  the  anode  to  the  cathode.  When  the  heat  generated 
by  the  passage  of  the  current  through  this  highly  resisting  medium  has 
melted  a  sufficient  amount  of  the  salt,  the  rod  is  removed,  and  the 
resistance  of  the  fused  material  suffices  to  maintain  the  temperature. 
The  calcium  rises  round  the  cathode  and  collects  on  the  surface  of 
the  bath.  By  slowly  elevating  the  copper  cathode,  the  calcium, 
which  adheres  to  it,  may  be  drawn  out  of  the  fused  mass  in  the  form 
of  a  gradually  lengthening,  irregular  rod.  The  rod  of  calcium  is 
kept  constantly  in  contact  with  the  metal  which  accumulates  on  the 
surface,  and  thus  forms  one  of  the  electrodes. 

Calcium  is  a  silver-white,  crystalline  metal  (m.-p.  750°,  density 
1.85)  which  is  a  little  harder  than  lead,  and  can  be  cut,  drawn,  and 
rolled.  Only  four  metals  (p.  645)  are  better  conductors  of  electricity. 
It  interacts  vigorously  with  water.  When  dry  and  cold  it  is  inactive, 
but  when  heated  it  unites  vigorously  with  hydrogen,  oxygen,  the  halo- 
gens, and  nitrogen.  It  burns  in  the  air,  giving  a  mixture  of  the  oxide 
and  nitride  Ca3N2.  The  presence  of  the  latter  may  be  shown  by  the 
liberation  of  ammonia  when  water  is  brought  in  contact  with  the 
residue: 

Ca3N2  +  6H20  ->  3Ca(OH)2  +  2NH3T. 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       703 

Calcium  Hydride  CaHz.  —  The  white  crystalline  hydride  CaH2 
is  formed  by  direct  union  of  the  constituents.  It  is  known  in 
commerce  as  hydrolyte,  and  is  an  expensive,  but  portable  source 
of  hydrogen  for  filling  war  balloons:  CaH2  +  2H2O  — >  Ca(OH)2 
+  2H2. 

Calcium  Chloride  CaCl2.  —  Calcium-ion  is  present  in  small 
amount  in  sea-water,  and  hence  compounds  containing  calcium 
chloride,  such  as  tachydrite  CaCl2,MgCl2,12H2O,  are  found  in  salt 
deposits.  The  salt,  for  which  there  is  no  extensive  commercial 
application  (cf.  p.  688),  is  formed  in  large  quantities  as  a  by-product 
in  several  industrial  operations.  Thus,  it  arises  in  the  liberation  of 
ammonia  from  ammonium  chloride  by  the  action  of  lime,  in  the 
manufacture  of  potassium  chlorate  (p.  481),  and  in  the  Solvay  soda 
process  (p.  688).  By  evaporation  of  any  solution  the  hexahydrate 
of  the  salt,  CaCl2,6H2O,  is  obtained  in  large,  deliquescent,  six-sided 
prisms.  On  account  of  the  great  concentration  of  a  saturated 
solution  of  this  compound,  the  solid  and  solution  do  not  reach  a 
condition  of  equilibrium  with  ice  (cf.  p.  200)  until  the  temperature  has 
fallen  to  —50°.  The  freezing  mixture  is  best  made  with  the  hydrate, 
and  not  with  the  anhydrous  salt,  as  the  latter  gives  out  much  heat 
in  becoming  hydrated  (dissolving) .  The  former  absorbs  heat  in  lique- 
fying, as  all  solids  do. 

There  are  several  other  hydrates  of  calcium  chloride  containing 
less  than  6H20,  and  those  containing  less  water  have  lower  aqueous 
tensions  (cf.  p.  153)  than  have  those  containing  more.  By  elevating 
the  temperature,  however,  it  is  easy  to  raise  the  aqueous  tension  even 
of  the  monohydrate  until  it  exceeds  the  partial  pressure  of  water  vapor 
in  ordinary  moist  air,  and  so  to  drive  out  the  water.  To  perform  this 
rapidly,  a  temperature  of  over  200°  is  required.  The  partially  de- 
hydrated calcium  chloride  CaCl2,2H2O  forms  a  porous  mass  which  is 
used  in  chemical  laboratories  for  drying  gases  and  liquids.  Usually 
the  dehydration  is  left  incomplete,  as,  at  the  temperature  required  to 
complete  it  rapidly,  some  interaction  with  the  water  occurs:  CaCl2  + 
H2O  — >  CaO  +  2HC1,  and  a  little  free  alkali  is  present  in  the  product. 
When  calcium  chloride  is  used  as  a  drying  agent,  it  is  naturally  able 
to  reduce  the  partial  pressure  of  the  water  vapor  only  to  the  value  of 
the  aqueous  tension  of  the  hydrate  which  is  present,  and  no  further. 
Even  at  low  temperatures  the  aqueous  tensions  of  hydrates  are 
always  perceptible  (cf.  p.  152).  Concentrated  sulphuric  acid  is  a 
more  thorough  drying  agent  than  is  calcium  chloride,  and  phosphorus 


704  INORGANIC  CHEMISTRY 

pentoxide,  whose  hydrated  form  (metaphosphoric  acid)  has  no  observ- 
able aqueous  tension,  is  better  still. 

Calcium  chloride  forms  molecular  compounds,  not  only  with 
water,  but  also  with  ammonia  CaCl2,8NH3  and  with  alcohol.  For 
drying  these  substances,  therefore,  quicklime  is  employed.  Hydro- 
gen sulphide  interacts  with  the  salt,  giving  hydrogen  chloride,  which 
renders  the  gas  impure.  This  gas  is  therefore  dried  with  phosphorus 
pentoxide. 

Calcium  Fluoride  CoF2.  —  This  compound  occurs  in  nature  as 
fluorite  or  fluor-spar  CaF2.  It  crystallizes  in  cubes,  is  insoluble  in 
water,  and  when  pure  is  colorless.  Natural  specimens  often  possess  a 
green  tint  or  show  a  violet  fluorescence.  It  is  formed  as  a  precipi- 
tate when  a  soluble  fluoride  is  added  to  a  solution  of  a  salt  of  calcium. 

Fluorite  is  used  in  the  etching  of  glass,  as  the  source  of  the  hydro- 
gen fluoride  (p.  281).  It  is  easily  fusible,  as  its  name  indicates  (Lat. 
fluere,  to  flow),  and  is  employed  in  metallurgical  operations,  for  the 
purpose  of  lowering  the  melting-point  (or  freezing-point,  which  is 
the  same  thing,  cf.  p.  144)  of  the  slag  (p.  653),  and  so  facilitating 
the  separation  of  the  latter  from  the  metal. 

Calcium  Carbonate  CaCOs.  —  This  compound  is  found  very 
plentifully  in  nature.  Limestone  is  a  compact,  indistinctly  crystal- 
line variety,  while  marble  is  a  distinctly  crystalline  form.  Chalk* 
is  a  deposit  consisting  of  the  calcareous  parts  of  minute  organisms; 
and  egg-shells,  oyster-shells,  coral,  and  pearls  are  other  varieties  of 
organic  origin. f  A  laminated  kind  of  limestone  found  at  Solnhofen 
is  used  for  lithographic  work.  Calcite  and  Iceland  spar  (Ger.  spalten, 
to  split)  are  pure  crystallized  calcium  carbonate.  The  former  occurs 
in  flat  rhombohedrons,  or  in  pointed,  six-sided  crystals  (Fig.  68,  p.  172) 
known  as  scalenohedrons  (" dog-tooth"  spar)  belonging  to  the  same 
system.  All  the  crystals  split  with  ease  parallel  to  three  planes  of 
cleavage,  giving  rhombohedrons  of  the  shape  shown  in  Fig.  69  (p.  173), 
but  this  nearly  cubical  form  is  itself  seldom  found  in  nature.  An 
entirely  different  crystallized  variety  is  known  as  aragonite.  This 
belongs  to  the  rhombic  system,  although  complex  crystals  ("twins") 
of  hexagonal  outline  constitute  the  most  familiar  specimens.  Ara- 
gonite, when  heated  gently,  resolves  itself  into  a  mass  of  minute 

*  Blackboard  "crayon"  is  usually  made  of  gypsum,  and  not  of  chalk, 
t  The  hard  coverings  of  Crustacea  and  insects  are  not  made  of  this  substance, 
but  of  an  organic  material  called  chit  in. 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      705 

crystals  of  calcite,  and  the  latter  is  the  more  stable  form  of  the  sub- 
stance. When  calcium  carbonate  is  produced  by  precipitation  it  is 
at  first  amorphous  but  slowly  becomes  crystalline.  In  cold  liquids 
the  resulting  crystals  are  calcite;  but  in  warm  solutions  the  less  stable 
form,  that  of  aragonite,  is  first  assumed. 

When  heated,  calcium  carbonate  dissociates,  giving  carbon  diox- 
ide and  quicklime 

CaC03  <=*  CaO  +  C02. 

At  ordinary  temperatures  the  decomposition  is  imperceptible.  On 
the  contrary,  atmospheric  carbon  dioxide,  in  spite  of  its  very  low 
partial  pressure,  combines  with  quicklime,  giving  " air-slaked"  lime. 
As  the  temperature  rises,  however,  the  tension  of  carbon  dioxide 
coming  from  the  carbonate  increases,  and  has  a  fixed  value  for  each 
temperature.  If  it  is  continually  allowed  to  escape,  so  that  the 
maximum  pressure  is  not  reached,  the  whole  of  the  salt  eventually 
decomposes.  If,  on  the  other  hand,  the  gas  is  confined,  the  system 
reaches  a  condition  of  equilibrium.  Attempts  to  increase  the  pres- 
sure on  the  gas  beyond  the  dissociation  pressure  proper  to  the  existing 
temperature,  result  in  recombination.  At  700°  the  pressure  is  only  25 
mm.,  at  900°  it  reaches  an  atmosphere,  and  at  950°  two  atmospheres. 
The  phenomenon  is  similar  to  the  dissociation  of  barium  peroxide 
(p.  82)  and  to  the  evaporation  of  a  liquid  (p.  145). 

Limestone  is  used  in  the  manufacture  of  quicklime  (#.#.)  and  of 
glass.  It  is  employed  largely  as  a  flux  in  metallurgy,  when  minerals 
rich  in  silica  are  brought  into  fusible  form  by  the  production  of  calcium 
silicate  CaSiO3.  Large  amounts  also  find  application  as  building- 
stone. 

Limestone  is  soluble  in  water  containing  carbonic  acid,  giving 
calcium  bicarbonate  Ca(HCO3)2  (p.  141,  also  see  p.  722).  By 
solution  of  limestone,  caves  are  often  formed.  Conversely,  sub- 
terranean water  containing  the  bicarbonate,  when  it  reaches  such  a 
cavern,  loses  carbon  dioxide  and  deposits  calcium  carbonate  as 
stalactites  or  columns  hanging  from  the  ceiling.  The  drippings 
form  stalagmites  on  the  floors. 

The  Phase  Rule,  a  Method  of  Classifying  all  Systems  in 
Equilibrium.  —  The  formal  resemblance  that  we  have  just  shown 
to  exist  between  the  modes  of  behavior  of  a  system  composed  of 
water  and  water-vapor  in  physical  equilibrium,  on  the  one  hand,  and 
of  a  system  made  up  of  calcium  oxide,  carbon  dioxide,  and  undecom- 


706  INORGANIC  CHEMISTRY 

posed  calcium  carbonate  in  chemical  equilibrium,  on  the  other,  is  not  a 
coincidence.  A  study  of  all  kinds  of  systems  in  equilibrium  shows 
that  their  different  modes  of  behavior  are  limited  in  variety  and  can 
be  classified  in  a  very  simple  way. 

The  categories  used  for  classification  are:  (1)  the  independent 
components  in  the  system,  and  their  number ;  (2)  the  distinct,  physi- 
cally separable  parts  or  phases  (p.  192)  of  the  system,  and  their 
number;  and  (3)  the  conditions  —  temperature,  pressure,  and  concen- 
tration; and  (4)  the  degree  of  variability  in  the  conditions  which  is  pos- 
sible without  the  occurrence  of  a  change  in  the  number  of  the  phases. 

The  mode  of  employment  of  these  three  categories  may  be  illus- 
trated in  the  order  of  their  mention: 

1.  In  the  water  and  water-vapor  system,  water  is  the  only  com- 
ponent.    In  the  calcium  carbonate  system,  the  independent  com- 
ponents are  two  in  number,  calcium  oxide  and  carbon  dioxide. 

2.  In  the  water  and  water-vapor  system  there  are  two  phases,  the 
liquid  phase  and  the  vapor  phase.     In  the  calcium  carbonate  system 
there  are  three  phases  —  two  solid  phases,  the  carbonate  and  oxide, 
and  one  gaseous  phase,  the  carbon  dioxide. 

3.  In  the  water  and  water-vapor  system  either  the  temperature 
or  the  pressure  may  be  altered,  within  certain  limits,  at  will'.     But, 
whichever  one  of  these  two  conditions  it  be  that  is  thus  changed,  the 
preservation  of  the  two  phases  will  at  once  require  a  simultaneous 
modification  in  the  other  condition,  of  such  a  nature  as  will  suit  the 
new  value  of  the  first.     Thus,  if  the  pressure  upon  the  vapor  is  raised, 
the  vapor  phase  will  be  destroyed  (p.  146)  unless  the  temperature  is 
simultaneously  elevated  to  a  certain  definite  point  (p.  146).     Sim- 
ilarly, if  the  temperature  is  raised,  the  liquid  phase  will  pass  into 
vapor  unless  a  sufficient  increase  in  the  pressure  is  simultaneously 
effected.     There  is  therefore  one,  and  only  one  degree  of  variability  in 
the  conditions  —  the  system  is  univariant.     The  experimenter  has 
only  one  free  choice,  and  after  he  has  made  this  choice,  the  other 
condition  necessary  to  preserve  the  number  of  phases  is  determined 
by  the  system.     By  a  study  of  the  calcium  carbonate  system,  as 
described  above,  it  will  be  seen  that  it  also  is  a  univariant  system. 

A  partial  generalization  of  these  results  leads  to  the  conclusion 
that  when  the  number  of  the  phases  exceeds  the  number  of  the  compo- 
nents by  one,  the  system  is  univariant.  Additional  illustrations  are 
now  required  for  reaching  a  still  more  general  statement. 

If  ice  be  added  to  the  water  and  water-vapor  system,  and  the 
system  be  allowed  to  reach  equilibrium  with  all  three  phases  present, 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       707 

we  find  on  analyzing  as  before:  one  component  (water),  three  phases 
(solid,  liquid,  and  gaseous),  and  no  variability  in  the  system.  Neither 
temperature  nor  pressure  may  be  altered  without  ensuing  disappear- 
ance of  one  or  other  of  the  three  phases.  This  system  is  therefore 
invariant. 

It  thus  appears  that  with  an  equal  number  of  components,  the 
more  phases  we  have,  the  more  restricted  are  the  possibilities  of 
change  in  the  conditions.  When  the  number  of  phases  equals  the 
number  of  components,  the  system  is  bi variant;  when  the  number  of 
phases  exceeds  the  number  of  components  by  one,  the  system  is  uni- 
variant;  when  the  number  of  phases  exceeds  the  number  of  compo- 
nents by  two,  the  system  is  invariant,  and,  in  general, 

Components  +  2  =  Phases  +  Variable  conditions. 

The  law,  of  which  this  equation  is  the  most  compact  expression,  is 
known  as  the  phase  rule,  and  was  first  formulated  by  Willard  Gibbs, 
of  Yale  University.  It  applies  to  physical  and  chemical  equilibria 
without  distinction,  and  involves  no  consideration  of  molecules  or 
atoms. 

Thus,  the  formal  resemblance  between  the  dissociation  phenomena 
exhibited  by  calcium  carbonate  and  other  compounds,  on  the  one  hand, 
and  the  behavior  of  a  liquid  in  contact  with  its  vapor  on  the  other,  is 
due  simply  to  the  fact  that  in  each  case  the  number  of  phases  exceeds 
the  number  of  components  by  one.  This  will  be  found  to  hold  in  all 
cases  where  there  is,  at  each  temperature,  a  constant  dissociation 
pressure.  A  decomposing  hydrate,  for  example,  furnishes  such  a 
case.  The  system  is  made  up  of  one  gaseous  phase  (water-vapor) 
and  two  solid  phases  (the  hydrate,  and  the  anhydrous  substance  or  a 
lower  hydrate).  It  has  three  phases  and  two  components  (water  and 
the  anhydrous  substance),  and  is,  therefore,  uni variant. 

Again,  when  we  have  a  sharp  transition-point  at  a  fixed  temper- 
ature, that  is,  a  unique  temperature  at  which  alone  several  different 
states  of  aggregation  of  a  substance  can  co-exist  (p.  144),  the  system 
is  always  invariant.  Thus,  ice  and  water  (and  vapor)  co-exist  at  the 
melting-point  of  ice:  three  phases  and  one  component.  Again,  at 
96°  two  solid  forms  of  sulphur  (p.  412)  co-exist  with  sulphur  vapor. 
In  these  cases  the  change  which  takes  place  at  the  transition  point  is 
purely  physical.  Analogous  cases  in  which  the  change  is  a  chemical 
one  are  equally  familiar.  The  decahydrate  of  sodium  carbonate 
decomposes  (p.  689)  above  35.2°.  At  this  temperature  the  deca- 
hydrate and  the  monohydrate  co-exist  with  the  saturated  solution 


708 


INORGANIC  CHEMISTRY 


and  water-vapor:  four  phases  and  two  components.  The  system  is, 
therefore,  invariant.  The  cases  of  gypsum  (see  p.  718)  and  sodium 
sulphate  (p.  193)  are  similar. 

In  the  case  of  barium  peroxide,  it  has  been  found  by  Hildebrand 
that,  as  a  given  specimen  decomposes  at  a  fixed  temperature,  it  does 
not  give  a  constant  pressure  of  oxygen,  but  a  continuously  diminish- 
ing pressure.  This  shows  that  the  two  solid  phases,  BaO  and  Ba02, 
are  not  independent  of  one  another,  otherwise  the  latter  would  give 
the  same  pressure  of  oxygen  until  it  was  all  decomposed.  These 
substances  must  be  soluble  in  one  another,  giving  a  solid  solution, 
and  therefore  a  single  phase.  This  introduces  a  second  variable, 
namely  the  concentration  of  this  solid  solution.  The  same  thing  has 
been  observed  with  the  carbonates,  so  that  CaO  and  CaCOs  are 
mutually  soluble,  and  the  dissociation  pressure  is  not  strictly  constant, 
but  varies  with  the  proportions  in  which  the  solids  are  present,  as 
well  as  with  the  temperature. 

Calcium  Oxide  CaO.  —  Pure  oxide  of  calcium  (quicklime)  may 
be  made  by  ignition  of  pure  marble  or  calcite.  For  commercial 
purposes  limestone  is  converted  into  quicklime  in  kilns  (Fig.  145). 

In  the  United  States  the  "  long-flame" 
process,  in  which  the  kiln  is  first 
charged  with  limestone  and  a  fire 
is  then  kindled  in  a  cavity  left  at 
the  bottom,  is  the  one  most  com- 
monly used.  Elsewhere,  the  lime- 
stone and  coal  are  thrown,  in  alter- 
nate layers,  into  the  kiln,  and  the 
products  are  withdrawn  at  the  bot- 
tom. The  latter,  the  " short-flame" 
method,  demands  less  fuel,  since  the 
operation  is  continuous,  and  the 
structure  is  never  allowed  to  cool, 
but  the  quicklime  is  mixed  with  the 
ash  of  the  coal.  In  both  cases,  the 
flames  and  heated  gases  from  the  fire 

pass  between  the  pieces  of  limestone,  and  the  carbon  dioxide  lib- 
erated is  carried  off  by  the  draft.  When  the  gas  is  to  be  used 
in  the  Solvay  process  or  in  the  refining  of  sugar,  coke  (smokeless), 
instead  of  coal,  is  employed  as  the  fuel.  As  low  a  temperature 
as  possible  is  used.  A  high  temperature  causes  impurities  in  the 


CARBON 
""*"  oioxioe 


FIG.  145. 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       709 

limestone  (e.g.,  clay)  to  interact  with  the  quicklime,  giving  fusible 
silicates,  which  fill  the  pores  and  interfere  with  the  subsequent  slaking 
with  water.  Since  the  change  is  reversible,  if  the  gas  lingers  in  the 
kiln  (at  760  mm.  pressure),  a  temperature  over  900°  is  required 
to  drive  the  action  forward  (p.  705).  Hence,  a  good  draft,  which 
removes  the  gas  as  fast  as  it  is  formed,  permits  the  use  of  a  lower 
temperature. 

Pure  calcium  oxide  is  a  white,  porous  solid.  It  is  infusible  even  in 
the  oxyhydrogen  flame,  but  may  be  melted  and  boiled  in  the  electric 
arc.  It  is  not  reducible  by  sodium,  or  by  carbon  excepting  at  the 
temperature  of  the  electric  furnace. 

Calcium  Hydroxide  Ca(Ofl")2.  —  When  water  is  poured  upon 
quicklime  it  is  first  absorbed  into  the  pores  mechanically.  The 
chemical  union  by  which  the  hydroxide  is  formed: 

CaO  +  H20  <=*  Ca(OH)2, 

proceeds  slowly.  When  it  is  complete,  the  product  is  a  fine  powder 
occupying  a  larger  volume  than  the  original  materials.  Much 
heat  is  evolved,  and  part  of  the  water  is  driven  off  as  steam.  The 
action  is  reversible,  and  at  a  high  temperature  the  hydroxide  can  be 
dehydrated:  Quicklime  from  pure  limestone  slakes  easily,  and  is 
known  as  "fat"  lime.  That  made  from  material  containing  clay  or 
magnesium  carbonate  is  "poor"  lime.  The  latter  slakes  slowly 
and  often  incompletely,  and,  when  used  for  mortar,  does  not  harden 
so  satisfactorily. 

Calcium  hydroxide  is  slightly  soluble  in  water  (limewater),  600 
parts  of  water  dissolving  1  part  of  the  hydroxide  at  18°,  about  twice 
as  much  water  being  required  at  100°.  The  solution,  relatively  to 
its  concentration,  is  strongly  alkaline.  On  account  of  its  cheapness, 
this  substance  is  used  by  manufacturers  in  almost  all  operations 
requiring  a  base,  and  it  thus  occupies  the  same  position  amongst 
bases  that  sulphuric  acid  does  amongst  acids.  When  the  presence  of 
much  water  is  unnecessary  or  undesirable,  a  suspension  of  the  solid 
hydroxide  in  the  saturated  solution  ("milk  of  lime"),  or  even  a  paste, 
is  employed.  In  such  cases,  as  in  the  manufacture  of  caustic  alkalies 
(p.  666),  the  action  takes  place  with  the  part  which  is  at  the  moment 
in  solution,  and  proceeds  through  the  continual  readjustment  of  a 
complex  set  of  equilibria.  Caustic  lime  is  employed  in  the  manu- 
facture of  alkalies  (p.  666),  bleaching  powder,  and  mortar  (see 


710  INORGANIC  CHEMISTRY 

below),  the  removal  of  the  hair  from  hides  in  preparation  for  tanning, 
the  softening  of  water,  and  as  a  whitewash. 

Mortar.  —  Mortar  is  made  by  mixing  water  with  slaked  lime 
and  a  large  proportion  of  sand.  The  " hardening"  process  consists 
in  an  interaction  of  the  carbon  dioxide  of  the  air  with  the  calcium 
hydroxide : 

C02  +  Ca(OH)2  -»  CaC03  +  H2O. 

After  the  superficial  parts  have  been  changed,  the  process  goes  on 
very  slowly,  and  many  years  are  required  before  the  deeper  layers 
have  been  transformed.  The  minute  crystals  of  calcium  carbonate 
which  are  formed  are  interlaced  with  the  sand  particles,  and  a  rigid 
mass  is  finally  produced.  The  sand  is  useful  in  two  ways.  In  the 
first  place,  it  makes  the  whole  material  more  porous,  and  so  facilitates 
the  diffusion  of  the  gas  into  the-  interior.  In  the  second  place,  since 
the  sand  is  not  itself  altered,  its  presence  prevents  the  formation  of 
large  cracks  which  would  otherwise  arise  from  the  shrinkage  that 
accompanies  the  formation  of  the  carbonate.  The  " hardening" 
does  not  begin  until  the  excess  of  water  used  in  making  the  mortar  has 
evaporated,  and  hence  ordinary  mortar  is  unsuitable  for  use  in  damp 
places  such  as  cellars. 

Calcium  Oxalate  CaC2O^.  —  This  salt  may  be  observed  under 
the  microscope  in  the  cells  of  many  plants.  It  appears  in  the  form 
of  needle-shaped  or  of  granular  crystals.  Since  it  is  the  least  soluble 
salt  of  calcium,  its  formation  is  used  as  a  test  for  calcium-ion.  Cal- 
cium is  estimated  quantitatively  by  adding  ammonium  oxalate  to 
the  neutral  or  slightly  alkaline  solution  of  the  calcium  salt.  The 
precipitate  is  separated  by  filtration,  washed  with  water,  and  then 
heated  strongly  (ignited)  in  a  crucible.  The  product  weighed  is 
calcium  oxide,  CaC204  — •»  CaO  +  CO2  +  CO.  More  often,  perhaps, 
the  oxalate  is  ignited  with  sulphuric  acid,  and  the  calcium  weighed  as 
sulphate. 

Theory  of  Precipitation.  —  The  precipitation  of  calcium  oxa- 
late CaC2O4,  just  referred  to,  is  a  typical  one  and  may  be  used  to 
illustrate  the  application  of  ion-product  constancy  (p.  698)  to  explain- 
ing the  phenomenon.  The  same  explanation  serves  for  all  precipita- 
tions involving  double  decomposition. 

The  first  thing  to  be  remembered  is  that  the  precipitate  which  we 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      711 

observe,  however  insoluble  its  material  may  be,  does  not  include  all 
of  the  substance,  but  only  the  excess  beyond  what  is  required  to 
saturate  the  water.  The  liquid  surrounding  the  precipitate  is  always 
a  saturated  solution  of  the  substance  precipitated.  If  it  were  not  so, 
some  of  the  precipitate  would  dissolve  until  the  liquid  became  satu- 
rated. Thus,  for  example,  when  we  add  ammonium  oxalate  solution 
to  calcium  chloride  solution: 


~C^04    **  CaC2°4  (dslvd)  **  CaC2°4  (solid)' 


2C 

the  liquid  is  a  saturated  solution  of  calcium  oxalate,  with  the  excess 
of  this  salt  susp'ended  in  it  as  a  precipitate. 

Looking  at  the  matter  from  this  viewpoint,  we  perceive  the 
application  of  the  rule  of  ion-product  constancy.  In  this  saturated 
solution  (p.  698)  the  product  of  the  ion-concentrations,  [Ca++]  X 
[C2O4=],  is  constant.  If  the  original  solutions  had  been  so  very  dilute 
that,  when  they  were  mixed,  the  product  of  the  concentrations  of 
these  two  ions  had  not  reached  the  value  of  this  constant,  no  precipita- 
tion would  have  occurred.  As  a  matter  of  fact  the  ion-product  consid- 
erably exceeded  the  requisite  value,  and  hence  the  salt  was  thrown 
down  until  the  balance  remaining  gave  the  numerical  value  in  ques- 
tion. The  rule  for  precipitation,  then,  is  as  follows:  Whenever  the 
product  of  the  concentrations  of  any  two  ions  in  a  mixture  exceeds  the 
value  of  the  ion-product  in  a  saturated  solution  of  the  compound 
formed  by  their  union,  this  compound  will  be  precipitated.  Naturally 
the  substances  with  small  solubilities,  and  therefore  small  ion-prod- 
uct constants,  are  the  ones  most  frequently  formed  as  precipitates. 

In  the  case  of  calcium  oxalate,  the  molar  solubility  (see  Table) 
is  0.0443.  In  so  dilute  a  solution  the  substance,  being  a  salt  (p. 
369),  must  be  practically  all  ionized.  Each  molecule  gives  one 
ion  of  each  kind.  The  molar  concentrations  of  these  ionic  sub- 
stances, Ca"1"1"  and  C2O4=,  in  the  solution,  when  the  solid  is  also  present, 
must  therefore  be  (practically)  0.0443,  each.  The  product  [Ca++] 
X  [C2O4=:]  is  thus  equal  to  0.0443  X  0.0443  or  0.08185.  If  in  mixing 
the  solutions,  exactly  equivalent  quantities  were  not  employed,  the 
values  of  the  two  factors  will  not  be  equal,  but  the  product  will  in  any 
case  possess  this  value. 

Rule  for  Solution  of  Substances.  —  The  rule  for  solution  of  any 

ionogen  follows  at  once  from  the  foregoing  considerations,  and  may  be 
formulated  by  changing  a  few  of  the  words  in  the  rule  just  given: 


712  INORGANIC  CHEMISTRY 

Whenever  the  product  of  the  concentrations  of  any  two  ions  in  a  mix- 
ture is  less  than  the  value  of  the  ion-product  in  a  saturated  solution 
of  the  compound  formed  by  their  union,  this  compound,  if  present  in 
the  solid  form,  will  be  dissolved.  When  applied  to  the  simplest  case, 
this  rule  means  that  a  substance  will  dissolve  in  a  liquid  not  yet  satu- 
rated with  it,  but  will  not  dissolve  in  a  liquid  already  saturated  with 
the  same  material.  The  value  of  the  rule  lies  in  its  application  to  the 
less  simple,  but  equally  common  cases,  such  as  when  an  insolu- 
ble body  is  dissolved  by  interaction  with  another  substance  (next 
section). 

Applications  of  the  Rule  for  Solution  to  the  Solution  of 
Insoluble  Substances.  —  So  long  as  a  substance  remains  in  pure 
water  its  solubility  is  fixed.  Thus,  with  calcium  hydroxide  at  18°, 
the  system  comes  to  rest  when  0.17  g.  per  100  c.c.  of  water  (0.02  moles 
per  liter)  have  gone  into  solution : 

Ca(OH)2  (solid)  ?±  Ca(OH)2  (dslvd)  <=»  Ca++  +  2OH~ 

But  if  an  additional  reagent  which  can  combine  with  either  one  of  the 
ions  is  added,  the  concentration  of  this  ion  at  once  becomes  less,  the 
ion-product  therefore  tends  to  diminish,  and  further  solution  must 
take  place  to  restore  its  value.  Thus,  if  a  little  of  an  acid  (giving  H+) 
be  added  to  the  solution  of  calcium  hydroxide,  the  union  of  OH~  and 
H+  to  form  water  removes  the  OH~,  and  solution  of  the  hydroxide 
proceeds  until  the  acid  is  used  up.  There  are  now  more  Ca4"1"  than 
OH~  ions  present,  but  the  ion-product  reaches  the  same  value  as  before, 
and  then  the  change  ceases.  If  a  further  supply  of  acid  is  added,  the 
removal  of  OH"  to  form  H2O  begins  again.  With  excess  of  the  acid, 
the  only  stable  OH~  concentration  is  that  which  is  a  factor  in  the  very 
minute  ion-product  of  water,  [OH~]  X  [H+],  which  is  0.061  X  0.061, 
or  O.Oisl.  Hence,  with  excess  of  acid,  the  calcium  hydroxide,  which 
requires  in  general  a  much  higher  concentration  of  OH~  than  this  to 
precipitate  it  or  to  keep  it  out  of  solution,  finally  all  dissolves. 

More  specifically,  if  we  assume  that  the  calcium  hydroxide  is 
wholly  dissociated  in  so  dilute  a  solution  (which  is  nearly  true), 
each  molecule  forms  one  ion  of  Ca4"1"  and  two  ions  of  OH~.  That 
is,  each  mole  of  Ca(OH)2  gives  one  mole  of  Ca^  and  two  moles 
of  OH~.  As  the  saturated  solution  contains  0.02  moles  of  the 
base,  the  molar  concentration  (assuming  complete  dissociation) 
of  [Ca++]  is  0.02  and  of  [OH~]  is  0.04.  Now,  the  ion-product  is 
the  product  of  the  concentrations  of  all  the  ions  formed,  i.e.,  Ca++, 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      713 

OH~,  and  OH~.  The  value  of  the  product  is  therefore  [Ca++] 
X  [OH-]  X  [OH-]  or  [Ca++]  X  [OH~]2.  That  is,  0.02  X  0.042  - 
0.0432.  Note  that  if  the  molecule  gives  two  (or  three)  ions  of  the 
same  kind,  the  whole  concentration  of  that  ion  taken,  and  is  also 
raised  to  the  second  (or  third)  power. 

This  particular  action  is  a  neutralization  of  an  insoluble  base. 
But  the  other  kinds  of  actions  by  which  insoluble  ionogens  pass  into 
solution  all  resemble  it  closely,  and  differ  only  in  details.  The  gen- 
eral outlines  of  the  explanation  are  the  same  in  every  case.  We  pro- 
ceed now  to  apply  it  to  the  common  phenomenon  of  the  solution  of  an 
insoluble  salt  by  an  acid. 

Interaction  of  Insoluble  Salts  with  Acids,  Resulting  in 
Solution  of  the  Salt.  —  Calcium  oxalate  passes  into  solution  when 
in  contact  with  acids,  especially  active  acids.  With  hydrochloric 
acid,  it  gives  calcium  chloride  and  oxalic  acid,  both  of  which  are 
soluble: 

CaC2O4  T  +  2HC1  ^  CaCl2  +  H2C2O4.  (1) 

The  action  of  acids  upon  insoluble  salts  is  so  frequently  mentioned  in 
chemistry  and  is  so  important  a  factor  in  analytical  operations  that  it 
demands  separate  discussion.  This  example  is  a  typical  one  and  may 
be  used  as  an  illustration. 

According  to  the  rules  already  explained  (p.  711),  calcium  oxalate 
(or  any  other  salt)  is  precipitated  when  the  numerical  value  of  the 
product  of  the  concentrations  of  the  two  requisite  ions  [Ca4"1"]  X 
[C2O4~]  exceeds  the  value  of  the  ion-product  for  a  saturated  solution 
of  calcium  oxalate  in  pure  water,  that  is,  exceeds  0.0gl85  (p.  711). 
When,  on  the  contrary,  the  product  of  the  concentrations  of  the  two 
ions  falls  below  the  limiting  value,  a  condition  which  may  arise  from 
the  removal  in  some  way  either  of  the  Ca++  or  of  the  C2O4~  ions,  the 
undissociated  molecules  will  ionize,  and  the  solid  will  dissolve  to 
replace  them  until  the  ionic  concentrations  necessary  for  equilibrium 
with  the  molecules  have  been  restored  or  until  the  whole  of  the  solid 
present  is  consumed.  Here  the  oxalate-ion  from  the  calcium  oxalate 
combines  with  the  hydrogen-ion  of  the  acid  (usually  an  active  one) 
which  has  been  added,  and  forms  molecular  oxalic  acid: 

C204=  +  2H+  fc>  H2C204.  (2) 

Hence,  dissociation  of  the  dissolved  molecules  of  calcium  oxalate  pro- 
ceeds, being  no  longer  balanced  by  encounters  and  unions  of  the  now 


714  INORGANIC  CHEMISTRY 

depleted  ions,  and  this  dissociation  in  turn  leads  to  solution  of  other 
molecules  from  the  precipitate. 

It  will  be  seen  that  the  removal  of  the  ions  in  this  fashion  can 
result  in  considerable  solution  of  the  salt  only  when  the  acid  produced 
is  a  feebly  ionized  one.  Here,  to  be  specific,  the  concentration  of 
the  C2O4~  in  the  oxalic  acid  equilibrium  (2)  above  must  be  less  than 
that  of  the  same  ion  in  a  saturated  calcium  oxalate  solution.  Now 
oxalic  acid  does  not  belong  to  the  least  active  class  of  acids,  and  its 
pure  solution  contains  a  considerable  concentration  of  C204=.  There 
is,  however,  a  decisive  factor  in  the  situation  which  we  have  not  yet 
taken  into  account.  The  hydrochloric  acid  which  we  used  for 
dissolving  the  precipitate  is  a  very  highly  ionized  acid  and  gives  an 
enormously  greater  concentration  of  hydrogen-ion  than  does  oxalic 
acid.  Hence  the  hydrogen-ion  is  in  excess  in  equation  (2),  and  the 

[-0+12  \/  rr^  Q  —  i 

condition  for  equilibrium,  •  -  '    -  —  =  K,  will  be  satisfied  by 


a  correspondingly  small  concentration  of  C2O4~.  In  this  particular 
case,  therefore,  the  [C204~]  of  the  oxalic  acid  ts  less  than  that  given 
by  the  calcium  oxalate.  The  whole  change,  therefore,  depends  for 
its  accomplishment,  not  only  on  the  mere  presence  of  hydrogen-ion, 
but  on  the  repression  of  the  ionization  of  the  oxalic  acid  by  the  great 
excess  of  hydrogen-ion  furnished  by  the  active  acid  that  has  been  used. 
As  a  matter  of  fact,  we  find  that  a  weak  acid  like  acetic  acid  has 
scarcely  any  effect  upon  a  precipitate  of  calcium  oxalate.  An  acid 
stronger  than  oxalic  acid  must  be  employed.  The  whole  scheme  of 
the  equilibria  is  as  follows: 

CaC204  (solid)  fc?  CaC204  (dslvd)  ^  Ca++-fC204=)_^  „  p  Q  ,  ,  ,    ,, 
2HC1  t*2Cl-+2H+    r 

When  excess  of  an  acid  sufficiently  active  to  furnish  a  large  concentra- 
tion of  hydrogen-ion  is  employed,  the  last  equilibrium  is  then  driven 
forward  and  the  others  follow.  With  addition  of  a  weak  acid,  only  a 
slight  displacement  occurs,  and  the  system  comes  to  rest  again  when 
the  molecular  oxalic  acid  has  reached  a  sufficient  concentration. 

A  generalization  may  be  based  on  these  considerations:  an  insol- 
uble salt  of  a  given  acid  will  in  general  interact  and  dissolve  when 
treated  with  a  solution  containing  another  acid,  provided  that  the 
latter  acid  is  a  much  more  highly  ionized  (more  active)  one  than  the 
former  (see  below). 

But  even  active  acids  frequently  fail  to  bring  salts  of  weak  acids 
into  solution,  especially  when  the  weak  acid  is  itself  present  also. 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       715 

Here  the  cause  lies  in  the  fact  that  such  salts  are  less  soluble  than 
those  of  the  calcium  oxalate  type,  and  give  so  low  a  concentration  of 
the  negative  ion  that  the  utmost  repression  of  the  ionization  of  the 
corresponding  acid  does  not  give  a  lower  value  for  the  concentration 
of  this  ion  than  does  the  salt  itself.  Thus,  we  have  seen  (p.  421)  that 
even  hydrochloric  acid  (dilute)  will  not  dissolve  a  number  of  sul- 
phides. For  example,  in  the  case  of  cupric  sulphide  in  a  solution 
saturated  with  hydrogen  sulphide,  the  S~  factor  in  the  solubility 
product  [Cu++]  X  [S=]  remains  smaller  than  that  in  the  scheme  defin- 
ing the  hydrogen  sulphide  equilibrium  [H+]2  X  [S=]  even  when  the 
[S=]  factor  in  the  latter  is  diminished  in  consequence  of  great  addition 
of  hydrogen-ion.  In  this  case,  the  first  link  in  the  chain  of  equilibria  : 

CuS  (solid)  ±5  CuS  (dslvd)  t?  Cu++  +  S= 
2HC1 


tends  so  decidedly  backward  that  only  the  use  of  concentrated  acid 
will  increase  the  concentration  of  the  H+  to  an  extent  sufficient  to 
secure  even  a  slight  advance  of  the  whole  action.  We  must  add,  there- 
fore, to  the  above  rule:  provided  also  that  the  salt  is  not  one  of 
extreme  insolubility.  This  point  will  be  illustrated  more  fully  in 
connection  with  the  description  of  individual  sulphides  (see  under 
Cadmium). 

Illustrations  of  the  application  of  these  generalizations  are  count- 
less. Carbonic  acid  is  made  from  marble  (p.  573),  hydrogen  sulphide 
from  ferrous  sulphide  (p.  419),  hydrogen  peroxide  from  sodium  per- 
oxide (p.  315),  and  phosphoric  acid  from  calcium  phosphate  (p.  548). 
In  each  case  the  acid  employed  to  decompose  the  salt  is  more  active 
than  the  acid  to  be  liberated.  On  the  other  hand,  calcium  oxalate 
and  calcium  phosphate  (except  when  freshly  precipitated)  are  in- 
soluble in  acetic  acid  because  this  acid  is  weaker  than  are  oxalic 
and  phosphoric  acids.  We  have  thus  only  to  examine  the  list  of 
acids  showing  their  degrees  of  ionization  (p.  367)  in  order  to  be  able 
to  tell  which  salts,  if  insoluble  in  water,  will  be  dissolved  by  acids  and, 
in  general,  what  acids  will  be  sufficiently  active  in  each  case  for  the 
purpose.  In  chemical  analysis  we  discriminate  between  salts  soluble 
in  water,  those  soluble  in  acetic  acid  (the  insoluble  carbonates  and 
some  sulphides,  FeS  and  ZnS,  for  example),  those  requiring  active 
mineral  acids  for  their  solution  (calcium  oxalate  and  the  more  insoluble 
sulphides,  for  example),  and  those  insoluble  in  all  acids  (barium 
sulphate  and  other  insoluble  salts  of  active  acids). 

The  influence  of  solubility  is  shown  not  only  by  the  sulphides,  but 


716  INORGANIC  CHEMISTRY 

also,  for  example,  by  the  sulphates.  Thus,  barium  sulphate  is  not  ap- 
preciably dissolved  even  by  the  most  active  acids  (p.  440),  being  a 
salt  of  a  rather  highly  ionized  acid,  and  being  itself  very  insoluble. 
Yet  calcium  sulphate,  being  much  less  insoluble  (p.  701),  is  dissolved 
to  a  noticeable  extent  by  the  same  acids. 

Precipitation  of  Insoluble  Salts  in  Presence  of  Acids.  — 

The  converse  of  solution,  namely,  precipitation,  depends  upon  the 
same  conditions:  an  insoluble  salt  which  is  dissolved  by  a  given  acid 
cannot  be  formed  by  precipitation  in  the  presence  of  this  acid.  Thus, 
calcium  oxalate  can  be  precipitated  in  presence  of  acetic  acid,  but  not 
in  presence  of  active  mineral  acids  in  ordinary  concentrations. 
Cupric  sulphide  or  barium  sulphate  can  be  precipitated  in  presence 
of  any  acid,  but  ferrous  sulphide  and  calcium  carbonate  only  in  the 
absence  of  acids. 

From  this  it  does  not  follow  that  calcium  oxalate,  for  example, 
cannot  be  precipitated  if  once  an  active  acid  has  been  added  to  the 
mixture.  To  secure  precipitation,  all  that  is  necessary  is  to  remove 
the  excess  of  hydrogen-ion  which  is  repressing  the  ionization  of  the 
oxalic  acid.  This  can  be  done  by  adding  a  base,  which  removes  the 
H+,  or  even  by  adding  sodium  acetate.  The  acetate-ion  C2H3O2~ 
unites  with  the  H+  to  form  the  little  ionized  acetic  acid,  in  presence 
of  which  calcium  oxalate  can  be  precipitated. 

Calcium  Carbide  CaC2.  —  The  manufacture  of  this  compound 
has  been  described  (p.  571),  and  the  formation  of  acetylene  by  its 
interaction  with  water  has  already  been  discussed  (p.  592).  The 
substance  was  discovered  by  Wohler  in  1862,  was  first  prepared  by 
the  use  of  electrical  heating  by  Borchers  in  1891,  and  was  made 
on  a  large  scale  in  1892  by  Willson,  a  Canadian  engineer.  The 
world's  production  in  1912  was  300,000  tons. 

Bleaching  Powder  Ca(OCl)CL  —  This  substance  (cf.  p.  475)  is 
manufactured  by  conducting  chlorine  into  the  lowest  of  a  series  of 
6-8  revolving  cylinders,  while  slaked  lime  is  fed  into  the  uppermost 
cylinder  (counter-current  system,  see  p.  687).  When  the  trans- 
formation has  reached  the  limit  (it  is  never  complete),  some  lime- 
dust  is  blown  in  to  absorb  the  remainder  of  the  free  chlorine.  The 
action  is  represented  by  the  equation  already  given,  and  not  by  the 
following : 

2Ca(OH)2  +  2C12  -*  CaCl2  +  Ca(OCl)2  +  2H2O. 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      717 

While  pure  lime  should  thus  yield  a  product  containing  49  per  cent  of 
chlorine,  in  practice  the  proportion  is  always  less.  Good  bleaching 
powder  should  contain  36-37  per  cent  of  chlorine. 

That  bleaching  powder  is  a  mixed  salt  CaCl(ClO)  rather  than  an 
equi-molar  mixture  of  calcium  chloride  and  calcium  hypochlorite 
(equation  above),  which  would  have  the  same  composition,  is  proven 
by  the  facts  that  the  material  is  not  deliquescent  as  is  calcium  chlo- 
ride, and  that  calcium  chloride  cannot  be  dissolved  out  of  it  by 
alcohol. 

Bleaching  powder  is  somewhat  soluble  in  water,  and  in  solution 
the  ions  Ca"1"1",  Cl"~,  and  C1O~  are  all  present.  Addition  of  acids 
causes  the  formation  of  hydrochloric  and  hypochlorous  acids.  The 
oxidizing  and,  incidentally,  the  bleaching  properties  (p.  475)  of  the 
latter  are  characteristic  of  the  acidified  liquid.  Weak  acids  like 
carbonic  acid  displace  the  hypochlorous  acid  only  (cf.  p.  476),  and 
hence  the  dry  powder,  when  exposed  to  the  air,  has  the  odor  of  the 
latter  substance  rather  than  that  of  chlorine. 

The  substance  is  largely  used  by  bleachers  (cf.  p.  477),  and  as  a 
disinfectant  to  destroy  germs  of  putrefaction  and  disease. 

Calcium  Nitrate  Ca(NOz)z»  —  This  salt  is  found  in  the  soil  (p. 
525),  and  may  best  be  prepared  in  pure  form  by  treating  marble 
with  nitric  acid  and  allowing  the  product  to  crystallize  from  the 
solution.  Calcium  nitrate  forms  several  hydrates.  The  tetrahydrate 
Ca(NO3)2,4H2O,  which  forms  transparent  monoclinic  prisms,  is  the 
one  deposited  at  ordinary  temperatures.  The  anhydrous  salt  is  easily 
soluble  in  alcohol.  It  is  used  in  the  laboratory  for  drying  nitrogen 
peroxide.  When  heated  it  decomposes  (cf.  p.  701),  giving  first  oxygen 
and  the  nitrite  and  then  nitrogen  peroxide  and  quicklime. 

Calcium  Sulphate  CaSO*.  —  This  salt  is  found  in  large  quan- 
tities in  nature.  The  mineral  anhydrite  CaSO4  occurs  in  the  salt 
layers  (see  under  Manganous  sulphate).  It  contains  no  water  of 
crystallization,  and  its  crystals  belong  to  the  rhombic  system.  The 
dihydrate,  CaS04,2H2O,  is  more  plentiful.  In  granular  masses  it 
constitutes  alabaster.  When  perfectly  crystallized  (monoclinic, 
Fig.  72,  p.  173)  it  is  named  gypsum  or  selenite.  The  same  hydrate 
is  formed  by  precipitation  from  solutions.  Its  solubility  is  about 
'1  in  500  at  18°.  Its  solubility  varies  in  an  unusual  manner  with 
temperature,  increasing  slowly  to  38°  and  then  falling  off. 

When  its  temperature  is  raised,  the  dihydrate  quickly  shows  an 


718  INORGANIC  CHEMISTRY 

appreciable  aqueous  tension.  After  three-fourths  of  the  water  has 
escaped,  a  hemihydrate  (CaS04)2,H20  remains,  which  shows  a  much 
smaller  tension  of  water  vapor  (cf.  p.  153). 

The  transition  temperature  at  Y/hich  the  dihydrate  passes  sharply  into  the 
hemihydrate  is  107°.  It  corresponds  to  the  temperature  of  35.2°  at  which  the 
decahydrate  of  sodium  carbonate  turns  into  the  monohydrate  and  water.  At  107° 
both  of  the  hydrates  are  in  equilibrium  with  the  solution  (p.  689).  Naturally  this 
system  can  exist  only  in  a  tube  sealed  up  to  prevent  the  escape  of  the  water. 

Plaster  of  paris  2CaSO4,H2O  is  manufactured  by  heating  gypsum 
until  nearly  all  the  water  of  hydration  has  been  driven  out.  When 
it  is  mixed  with  water,  the  dihydrate  is  quickly  re-formed  and  a 
rigid  mass  is  produced.  If,  in  course  of  manufacture,  the  water  is 
all  removed,  or  the  temperature  is  allowed  to  rise  much  above  the 
most  favorable  one  (about  125°),  the  product  when  mixed  with  water 
does  not  set  quickly  and  is  said  to  be  "  dead-burnt. "  In  explanation 
of  this  it  should  be  noted  that  natural  anhydrite  combines  very 
slowly  with  water.  Apparently  good  plaster  of  paris  must  contain 
some  unchanged  particles  of  the  dihydrate  which  may  act  as  nuclei. 
They  fulfil  the  same  role  as  the  crystal  which  is  added  to  a  super- 
saturated solution  (p.  193),  without  which  crystallization  may  be 
long  delayed  or  may  even  fail  to  take  place.  Probably,  with  moderate 
heating,  the  product  is  a  mixture  of  the  dihydrate  and  the  hemi- 
hydrate with  anhydrous  salt,  while  the  more  rapid  decomposition 
at  higher  temperatures  destroys  all  of  the  first.  The  former  mix- 
ture must  be  an  unstable  system,  and  the  dihydrate  loses  water 
to  the  anhydrous  salt.  At  ordinary  temperatures,  however,  this 
transference  must  be  very  slow,  and  hence  the  property  of  setting 
is  not  lost  by  prolonged  storage. 

That  the  plaster  sets,  instead  of  forming  a  loose  mass  of  dihydrate, 
is  due  to  the  fact  that  the  anhydrous  salt  is  more  soluble  than  the 
dihydrate,  and  so  a  constant  solution  of  the  one  and  deposition  of  the 
other  goes  on  until  the  hydration  is  complete: 

2CaS04,H20  (solid)  ^2CaS04  (dslvd)  ^2CaSo4,2H2O  (solid)! 

This  process  results  in  the  formation  of  an  interlaced  and  coherent 
mass  of  minute  crystals. 

Plaster  of  paris  is  used  for  making  casts  and  in  surgery.  The 
setting  of  the  material  is  accompanied  by  a  slight  increase  in  volume, 
and  hence  a  very  sharp  reproduction  of  all  the  details  in  the  structure 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      719 

of  the  mold  is  obtained.  An  "ivory"  surface,  which  makes  washing 
practicable,  is  conferred  by  painting  the  cast  with  a  solution  of  paraffin 
or  stearine  in  petroleum  ether.  The  waxy  material,  left  by  evapora- 
tion of  the  volatile  hydrocarbons,  fills  the  pores  and  prevents  solution 
and  disintegration  of  the  substance  by  water.  Stucco  is  made  with 
plaster  of  paris  and  rubble,  and  is  mixed  with  a  solution  of  size  or 
glue  instead  of  water. 

Calcium  Sulphide  CaS.  —  This  compound  is  most  easily  made 
by  strongly  heating  pulverized  calcium  sulphate  and  charcoal.  The 
sulphate  is  reduced:  4C  +  CaSO4 ->  CaS  +  4CO.  Calcium  sul- 
phide is  meagerly  soluble  in  water,  but  is  nevertheless  slowly 
dissolved  in  consequence  of  its  decomposition  by  hydrolysis  into 
calcium  hydroxide  and  calcium  hydrosulphide  (cf.  p.  421).  It  is 
probable  that  the  action  would  be  less  nearly  complete  than  it  is  if 
the  reverse  action  were  not  weakened  by  the  precipitation  of  the  cal- 
cium hydroxide: 

2CaS  +  2H20  *=?  Ca(OH)2|  +  Ca(SH)2. 

Since  calcium  sulphide  is  thus  decomposed  by  water  it  cannot  be 
precipitated  from  aqueous  solution  by  adding  a  soluble  sulphide,  such 
as  ammonium  sulphide,  to  a  solution  of  a  salt  of  calcium.  Only  the 
soluble  hydrosulphide  can  be  formed. 

Ordinary  calcium  sulphide,  after  it  has  been  exposed  to  sunlight, 
usually  shines  in  the  dark.  Barium  sulphide  behaves  in  the  same  way. 
On  this  account  these  substances  are  used  in  making  luminous  paint. 
They  apparently  owe  this  behavior  to  the  presence  of  traces  of 
compounds  of  vanadium  and  bismuth,  for  the  purified  substances  are 
not  affected  in  the  same  fashion.  Since  alkalies  decompose  proteins, 
moist  calcium  sulphide  is  used  as  a  depilatory. 

Phosphates  of  Calcium.  —  The  tertiary  orthophosphate  of  cal- 
cium Ca3(PO4)2,  known  as  phosphorite,  is  found  in  many  localities, 
and  is  often  derived  from  the  bones  of  animals.  Guano  contains 
some  of  the  same  substance,  along  with  nitrogen,  either  as  organic 
compounds  or  as  niter  (p.  525).  Apatite  3Ca3(PO4)2,CaF2,  a  double 
salt  with  calcium  fluoride  (or  chloride)  is  a  common  mineral  and  fre- 
quent component  of  rocks.  The  orthophosphate  forms  about  83 
per  cent  of  bone-ash  (p.  548),  and  is  contained  also  in  the  ashes  of 
plants.  It  may  be  precipitated  by  adding  a  soluble  phosphate  to  a 
solution  of  a  salt  of  calcium. 


720  INORGANIC   CHEMISTRY 

Since  it  is  a  salt  of  a  weak  acid,  and  belongs  to  the  less  insoluble 
class  of  such  salts,  calcium  phosphate  is  dissolved  by  dilute  mineral 
acids  (cf.  p.  714),  the  ions  HPO4=  and  H2P04~  being  formed.  When 
a  base,  such  as  ammonium  hydroxide,  is  added  to  the  solution,  the 
calcium  phosphate  is  reprecipitated  (cf.  p.  716). 

Calcium  phosphate  is  chiefly  used  in  the  manufacture  of  phos- 
phorus and  phosphoric  acid  (p.  548),  and  as  a  fertilizer.  The  supply 
of  calcium  phosphate  in  the  soil  arises  from  the  decomposition  of  rocks 
containing  phosphates,  and  is  gradually  exhausted  by  the  removal 
of  crops.  Bone-ash  is  sometimes  used  to  make  up  the  deficiency. 
It  is  almost  insoluble  in  water,  however,  and,  although  somewhat 
less  insoluble  in  natural  water  containing  salts  like  sodium  chloride 
(cf.  p.  715),  is  brought  into  a  condition  for  absorption  by  the  plants 
rather  slowly.  The  "superphosphate"  (see  below)  is  much  more 
soluble. 

Primary  calcium  orthophosphate  (superphosphate)  is  manufac- 
tured in  large  quantities  from  phosphorite  by  the  action  of  sul- 
phuric acid.  The  unconcentrated,  "chamber  acid"  is  used  for  this 
purpose,  as  water  is  required  in  the  resulting  action.  The  amounts  of 
material  employed  correspond  to  the  equation: 

Ca3(P04)2  +  2H2S04  +  6H2O  ->  Ca(H2P04)2;2H20  -f  2CaSO4,2H2O. 

As  soon  as  mixture  has  been  effected,  the  action  proceeds  with  evolu- 
tion of  heat,  and  a  large  cake  of  the  two  hydrated  salts  remains.  This 
mixture,  after  being  broken  up,  diied,  and  packed  in  bags,  is  sold  as 
"superphosphate  of  lime."  The  primary  phosphate  which  it  con- 
tains is  soluble  in  water,  and  is  therefore  of  great  value  as  a  fertilizer. 

Calcium  Cyanamide  CaCN2.  —  Calcium  carbide  (p.  571),  when 
strongly  heated,  absorbs  nitrogen,  giving  a  mixture  of  calcium 
cyanamide  and  carbon: 

CaC2  +  N2  -^  CaCN2  +  C, 

which  is  sold  as  nitro-lime  for  use  as  a  fertilizer.  When  treated 
with  hot  water,  the  cyanamide  is  hydrolyzed  into  calcium  carbon- 
ate and  ammonia : 

CaCN2  +  3H2O  ->  CaC03  +  2NH3. 

In  the  soil  the  decomposition  may  not  be  so  simple,  but  combined 
nitrogen  is  furnished  in  a  form  that  can  be  absorbed  by  plants. 
At  Odda  (Norway)  the  carbide  is  pulverized  and  placed  in  a 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       721 

cylindrical  furnace  (Fig.  146)  holding  300-450  kg.  The  heat  (800- 
1000°)  is  supplied  by  the  passage  of  an  electric  current  through 
a  thin  carbon  rod.  The  nitrogen  is  obtained  by  the 
fractionation  of  liquid  air  and  final  removal  of  all 
oxygen  by  passage  over  heated  copper,  and  is  forced 
in  under  pressure.  After  thirty-six  hours,  nitrogen  is 
no  longer  absorbed,  and  the  charge  is  pulverized  when 
cold. 

Sodium  cyanide  NaNC   is  now  manufactured  by 
fusing  nitro-lime  with  sodium  carbonate: 

CaCN2  +  C  +  Na2CO3  -*  CaC03  +  2NaNC. 

The  cyanide  is  extracted  from  the  insoluble  calcium 
carbonate  with  water,  in  which  it  is  exceedingly  sol- 
uble.    Sodium  cyanide  has  now  displaced  potassium  cyanide  in  the 
extraction  of  gold  from  its  ores. 

Nutrition  and  Fertilization  of  Crops.  —  The  plant  con- 
structs its  cellulose,  starch,  or  sugar,  and  secures  the  carbon-part  of 
all  its  organic  contents  from  the  carbon  dioxide  of  the  air  (p.  579). 
The  water  (90-95  per  cent  of  the  total  weight  of  the  plant)  comes 
from  the  soil  and  brings  up  in  solution  the  other  elements  required. 
All  soils  are  able  to  supply  sufficient  magnesium,  calcium,  and  iron, 
as  bicarbonates.  But  the  soil  may  lack:  sulphur,  absorbed  as 
sulphates;  nitrogen,  absorbed  chiefly  as  nitrates,  but  occasionally 
as  salts  of  ammonium;  potassium,  as  sulphate,  chloride,  or  nitrate; 
and  phosphorus,  as  soluble  phosphates.  The  soil  may  be  originally 
deficient  in  one  or  more  of  these  necessary  plant  foods,  or  the  supply 
may  have  been  exhausted  by  repeated  cropping.  Every  crop 
permanently  removes  certain  quantities.  For  example,  in  the  case  of 
nitrogen,  which  is  required  to  form  proteins  that  enter  largely  into 
the  fruit  (i.e.,  usually,  the  edible  part),  each  crop  of  Indian  corn  (45 
bushels)  removes  63  pounds  per  acre,  a  crop  of  cabbage  (15  tons) 
removes  100  pounds  per  acre,  clover  hay  (2  tons)  82  pounds,  and 
wheat  (15  bushels)  31  pounds.  When  the  store  in  the  soil  becomes 
meager,  the  crops  become  poor,  and  finally  cost  more  for  labor  than 
they  are  worth. 

Thus,  crops  have  to  be  fed,  just  like  cattle.  Moreover,  the 
elements  must  be  furnished  in  soluble  form  (cf.  pp.  559,  580,  627). 
Fertilizers  containing  potassium  (pp.  662,  674)  and  phosphorus 
(p.  559)  must  be  used,  when  the  soil  is  deficient  in  these  elements. 


722     .  INORGANIC  CHEMISTRY 

The  nitrogen  fertilizers  we  have  mentioned  are  sodium  nitrate 
(p.  685),  calcium  nitrate  (p.  525),  ammonium  sulphate  (p.  516), 
guano  and  manure  (p.  525),  "tankage"  and  ground  bones  from 
slaughter  houses,  calcium  cyanamide  (p.  720),  and  finally  the 
nitrates  from  bacterial  decomposition  of  root  nodules  (p.  515). 
That  systematic  use  of  fertilizers  does  influence  the  crops  is  indi- 
cated by  the  results  of  cultivation  of  land  which,  but  for  fertili- 
zation, would  long  since  have  become  almost  valueless.  The 
wheat  crop  per  acre,  being  the  average  of  ten  successive  years  is: 
Denmark  40  bushels,  Great  Britain  33,  Germany  29,  United 
States  14. 

Hard  Water.  —  As  we  have  seen  (pp.  576,  705),  limestone  (solu- 
bility, 0.013  g.  per  liter),  magnesium  carbonate  (soFty  1  g.  per 
liter),  and  iron  carbonate,  although  very  insoluble,  are  acted  upon 
by  the  carbonic  acid  in  natural  waters,  giving  bicarbonates  which 
are  roughly  about  thirty  times  as  soluble.  When  the  water  is 
boiled,  the  actions  (p.  576)  are  reversed,  and  the  carbonates  are 
reprecipitated.  These  bicarbonates  constitute  temporary  hardness, 
and  their  decomposition  produces  " fur"  in  a  kettle  and  boiler  crust 
in  a  boiler. 

The  sulphates  of  calcium  (sol'ty  2  g.  per  liter)  and  of  magnesium 
(sol'ty  354  g.  per  1.)  are  also  commonly  found  in  natural  waters. 
These  salts  are  not  affected  by  mere  boiling  (as  distinct  from  evap- 
oration) and  so,  along  with  magnesium  carbonate  (1  g.  per  1.)  and 
calcium  carbonate  (0.013  g.  per  1.)  give  permanent  hardness  to  the 
water. 

Hardness  is  estimated  in  "degrees."  In  France,  and  com- 
monly in  the  laboratory,  1  part  of  CaC03  (or  its  equivalent  of  other 
salts)  per  100,000  (0.01  g.  per  liter)  constitutes  one  degree.  In 
the  United  States  one  degree  is  1  grain  per  gallon  of  58,333  grains 
(0.017  g.  per  1.).  In  Britain  one  degree  is  1  grain  per  gallon  of  70,000 
grains  (0.014  g.  per  1.).  Well  water,  originating  in  chalk  or  lime- 
stone formations,  may  have  37°  (Fr.)  or  more  of  hardness. 

Damage  Due  to  Hardness  in  Water.  —  When  hard  water  is 
continually  fed  into  a  steam  boiler  and  only  steam  comes  out, 
naturally  the  salts  accumulate  and  produce  in  time  a  heavy  boiler 
crust,  which  settles  on  the  tubes.  Being  a  poor  conductor  of  heat 
compared  with  iron,  this  crust,  if  one-fourth  of  an  inch  thick,  will 
increase  the  consumption  (and  cost)  of  fuel  by  50  per  cent.  In 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      723 

addition,  the  iron,  not  being  in  direct  contact  with  water,  is  heated 
to  a  higher  temperature,  and  may  even  become  red  hot.  It  thus 
oxidizes  more  quickly  on  the  outside,  and  displaces  hydrogen  from 
water  (or  steam)  on  the  inside  (p.  116),  thus  changing  on  both  sides 
gradually  into  the  brittle  magnetic  oxide  Fe3O4.  If  the  crust  is  not 
removed,  or  prevented  (see  below),  the  life  of  the  boiler  is  greatly 
shortened,  and  a  serious  explosion  may  even  occur. 

In  washing,  in  the  household  or  laundry,  much  soap  is  wasted 
before  the  necessary  lather  is  secured.  The  soap,  for  example, 
the  sodium  stearate  (p.  620),  gives  magnesium  and  calcium  stea- 
rates,  which  are  insoluble,  forming  a  curd: 

CaS04  +  2Na(C02C17H35)  -»  Ca(C02Ci7H35)2 1  +  NasSO* 

The  permanent  solution  of  soap,  required  for  washing,  does  not 
begin  to  be  formed  until  all  the  hardness  has  thus  been  precipi- 
tated. Hence,  according  to  the  equation,  with  1°  (U.  S.)  hard- 
ness, 100  gallons  (U.  S.)  of  water  should  use  up  0.075  pounds  of 
soap  (1°  Brit,  and  100  gal.  Brit.,  0.075  lb.).  In  point  of  fact,  how- 
ever, the  colloidal  calcium  salts  adsorb  and  carry  down  with  them 
more  than  an  equal  amount  of  undecomposed  soap.  Hence,  actual 
measurement  shows  that,  with  1°  (U.  S.  or  Brit.)  of  hardness,  100 
gallons  (U.  S.  or  Brit.)  of  water  really  destroy  0.17  pounds  of  soap. 
Thus,  with  35°,  no  less  than  6  pounds  of  soap  per  100  gallons  are 
wasted  before  the  part  of  the  soap  that  is  to  do  the  work  begins  to 
dissolve. 

Treatment  of  Hard  Water.  —  The  temporary  hardness  can  be 
removed  by  boiling  the  water,  or  using  some  preheating  arrange- 
ment in  connection  with  the  boiler  (stationary  engines  only). 

Temporary  hardness  is  commonly  removed,  on  a  large  scale,  by 
adding  slaked  lime  (made  into  milk  of  lime)  in  exactly  the  quantity 
shown  by  an  analysis  of  the  water  to  be  required,  and.  stirring  for 
a  considerable  time: 

Ca(HC03)2  +  Ca(OH)2  ->  2CaC03 1  +  2H20.  (1) 

The  bicarbonate  is  neutralized  and  all  the  lime  precipitated.  The 
latter  is  removed  by  filtration. 

Permanent  hardness  is  not  affected  by  slaked  lime,  but  is  pre- 
cipitated by  adding  sodium  carbonate  in  the  necessary  proportion: 

CaS04  +  NajjCOs  ->  CaC03 1  +  NaaSO*  (2) 


724  INOBGANIC  CHEMISTRY 

When  both  kinds  of  hardness  are  present,  crude  caustic  soda 
(sodium  hydroxide)  may  be  employed.  It  neutralizes  the  bicar- 
bonate, precipitating  CaC03: 

Ca(HC03)2  +  2NaOH  -»  CaCO3  i  +  NasCOg  +  2H20,       (3) 


and  giving  sodium  carbonate.  The  latter  then  acts  as  in  equa- 
tion (2). 

Instead  of  this,  the  treatments  indicated  in  equations  (1)  and 
(2)  may  be  applied  in  combination  (Porter-Clark  process).* 

In  the  permutite  process,  the  water  is  simply  filtered  through 
sodium  silico-aluminate  (permutite  NaP,  an  artificial  zeolite),  which 
is  supplied  in  the  form  of  a  coarse,  insoluble  sand.  The  calcium, 
etc.,  in  the  water  is  exchanged  for  sodium,  which  does  no  harm: 

Ca(HC03)2  +  2NaP  ->  2NaHC03  +  CaP2. 

After  twelve  hours'  use,  the  permutite  is  covered  with  10  per  cent 
salt  solution,  and  allowed  to  remain  for  the  other  twelve  hours  of 
the  day,  when  it  is  ready  for  employment  once  more: 

2NaCl  +  CaP2  -»  CaCl2  +  2NaP. 

Only  salt,  which  is  inexpensive,  is  consumed,  and  calcium  chloride 
solution  is  thrown  away.  Permutite  removes  magnesium,  iron, 
manganese,  and  other  elements  in  the  same  way.  The  life  of  a 
charge  is  said  to  be  over  twenty  years. 

Hard  Water  in  the  Laundry.  —  As  we  have  seen  (p.  620), 
soap  will  soften  water,  but  the  calcium  and  magnesium  salts  of  the 
soap  acids,  which  are  precipitated,  are  sticky,  and  soil  the  goods 
being  washed.  Other  substances  that  soften  water,  not  only  give 
non-adhesive  precipitates,  but  are  also  much  cheaper,  and  an  at- 
tempt is  generally  made  to  utilize  them.  The  use  of  slaked  lime  is 
impracticable  on  a  small  scale. 

Washing  soda  NajjCOsjlOH^O  is  added  to  precipitate  both  kinds 
of  hardness: 


Ca(HC03)2  +  Na^CO,  -»  CaC03  +  2NaHCO3, 
CaSO4  +  Na^COs  -»  CaC03  -f  Na2S04. 


*  So  far  as  the  hardness  is  due  to  magnesium  bicarbonate,  a  double  propor- 
tion of  lime  must  be  added  to  precipitate  the  magnesium  as  hydroxide  (sol'ty 
0.01  g.  per  1.),  because  the  carbonate  is  too  soluble  (1  g.  per  1.). 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       725 

The  small  amounts  of  salts  of  sodium  which  remain  in  the  water 
have  no  action  on  soap. 

Household  Ammonia  NH^OH  acts  like  sodium  hydroxide  (p.  724) : 

Ca(HCO3)2  +  2NH4OH  ->  CaC03  +  (NH4)2CO3  +  2H20, 
CaSO4  +  (NH4)2CO3  ->  CaCO3  +  (NH4)2S04, 

except  that  it  will  not  precipitate  magnesium-ion. 

Borax  Na2B407,10H2O  (p.  639)  is  hydrolyzed  and  the  sodium  hy- 
droxide in  its  solution  acts  as  already  (p.  724)  described. 

The  supposed  bleaching  or  whitening  action  of  borax  or  soda 
is  a  myth;  these  salts  prevent  staining  by  the  iron  in  the  water. 
They  simply  precipitate  the  iron,  present  as  Fe(HCO3)2,  which 
almost  all  waters  contain,  as  FeC03  before  the  goods  are  put  in. 
This  precipitate  is  easily  washed  out  in  rinsing.  The  palmitate, 
etc.,  of  iron,  however,  which  the  soap  itself  would  throw  down,  is 
sticky  and  adheres  to  the  cloth.  The  air  subsequently  oxidizes 
it  (see  p.  801)  and  gives  hydrated  ferric  oxide  (rust),  which  is 
brownish-red. 

It  is  evident  that,  properly  to  achieve  their  purpose,  the  soda 
and  borax  must  be  added,  must  be  completely  dissolved,  and  must 
be  allowed  to  produce  the  precipitation  of  FeC03,  CaCO3,  etc.,  all 
before  the  soap  (or  the  goods)  is  introduced.  If  the  soap  is  dissolved 
before  or  with  the  soda,  it  will  take  part  in  the  precipitation,  and 
give  sticky  particles  containing  the  iron  and  calcium  salts  of  the 
soap  acids. 

The  soda,  borax,  and  ammonia  do  not  themselves  remove  dirt 
—  that  is  done  by  the  dissolved  soap  (p.  623).  With  the  help  of 
rubbing,  however,  they  do  emulsify  and  remove  animal  or  vege- 
table oil  and  grease,  but  not  mineral  oil  (p.  625),  when  these  happen 
to  be  on  the  goods.  But  soap  alone  will  do  this  also,  and  remove 
mineral  oil  as  well. 

Washing  powders  are,  or  ought  to  be,  mainly  sodium  carbonate, 
mixed  with  more  or  less  pulverized  soap. 

Calcium  Metasilicate  CaSiO3. —  This  salt  forms  the  mineral 
wollastonite,  which  is  rather  scarce,  and  enters  into  the  composi- 
tion of  many  complex  minerals,  such  as  garnet,  mica,  and  the  2eolites. 
It  may  be  made  by  precipitation  from  a  solution  of  sodium  meta- 
silicate  (p.  634),  or  by  fusing  together  powdered  quartz  and  calcium 
carbonate: 

SiO2  +  CaC03  -*  CaSiO3  +  CO2. 


726  INORGANIC  CHEMISTRY 

Glass.  —  Common  glass  is  a  complex  silicate  of  sodium  and  cal- 
cium, or  a  homogeneous  mixture  of  the  silicates  of  these  metals  with 
silica.  It  has  a  composition  represented  approximately  by  the 
formula  Na2O,CaO,6Si02,  and  is  made  by  melting  together  sodium 
carbonate,  limestone,  and  pure  sand: 

Na2C03  -f  CaC03  +  6Si02  -»  Na20,CaO,6Si02  +  2C02. 

For  the  most  fusible  glass,  a  smaller  proportion  of  quartz  is  employed. 
This  variety  is  known  as  soda-glass,  or,  from  its  easy  fusibility,  as 
soft  glass.  First,  the  materials  are  heated  to  a  temperature  high 
enough  to  produce  chemical  action  without  bringing  about  complete 
melting.  This  permits  the  ready  escape  of  the  gas.  Then  the  tem- 
perature is  raised  to  about  1200°  until  fusion  is  complete  and  all  the 
bubbles  have  escaped.  Finally,  the  crucible  and  its  contents  are 
allowed  to  cool  to  700-800°  to  permit  the  latter  to  acquire  the  vis- 
cosity required  for  working. 

Plate-glass  is  made  by  rolling  the  material  into  large  sheets  and 
polishing  the  surfaces  until  they  are  plane.  Window-glass  is  prepared 
by  blowing  bulbs  of  long  cylindrical  shape,  and  ripping  them  down  one 
side  with  the  help  of  a  diamond.  The  resulting  curved  sheets  are  then 
placed  on  a  flat  surface  in  a  furnace  and  are  there  allowed  to  open 
out.  Beads  are  made,  chiefly  in  Venice,  by  cutting  narrow  tubes 
into  very  short  sections  and  rounding  the  sharp  edges  by  fire.  Or- 
dinary apparatus  is  made  of  soft  soda-glass,  and  hence  when  heated 
strongly  it  tends  to  soften  and  also  to  confer  a  strong  yellow  tint 
(cf.  p.  692)  on  the  flame.  In  all  cases  the  articles  are  annealed  by 
being  passed  slowly  through  a  special  furnace  in  which  their  tempera- 
ture is  lowered  very  gradually.  Glass  which  has  been  suddenly 
chilled  is  in  a  state  of  tension  and  breaks  easily  when  handled. 

Bottles  are  made  with  impure  materials,  and  owe  their  color 
chiefly  to  the  silicate  of  iron  which  they  contain.  In  making  cheap 
glass,  sodium  sulphate  is  substituted  for  the  more  expensive  carbonate. 
In  this  case  powdered  charcoal  or  coal  is  added  to  reduce  the  sulphate : 

2Na2S04  +  C  +  2Si02  ->  2Na2Si03  +  C02  +  2S02. 

Soft  glass  is  partially  dissolved  by  water.  When  powdered  glass 
is  shaken  with  water,  sodium  silicate  dissolves  in  amount  sufficient 
to  give  an  alkaline  reaction  with  phenolphthalem  (cf.  p.  143). 

Bohemian,  or  hard  glass,  is  much  more  difficult  to  fuse  than  soda- 
glass,  and  is  also  much  less  soluble  in  water.  It  is  made  by  sub- 
stituting potassium  carbonate  for  sodium  carbonate.  Specially 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      727 

insoluble  glass,  for  laboratory  use,  such  as  Jena  and  non-sol  glass, 
is  made  with  boric  anhydride  B2O3,  in  addition  to  silica,  and  some 
zinc  oxide,  so  that  it  contains  borates  as  well  as  silicates.  When 
lead  oxide  is  employed  instead  of  limestone,  a  soda-lead  glass,  known 
as  flint  glass,  is  produced.  This  has  a  high  specific  gravity,  and  a 
great  refracting  power  for  light,  and  is  employed  for  making  glass 
ornaments.  By  the  use  of  grinding  machinery,  cut  glass  is  made 
from  it.  Engraving  on  glass  is  done  with  the  sand  blast. 

Colored  glass  is  prepared  by  adding  small  amounts  of  various 
oxides  to  the  usual  materials.  The  oxides  combine  with  the  silica,  and 
produce  strongly  colored  silicates.  Thus,  cobalt  oxide  gives  a  blue, 
chromium  oxide  or  cupric  oxide  a  green,  and  uranium  a  yellow  glass. 
Cuprous  oxide,  with  a  reducing  agent,  and  compounds  of  gold,  give 
the  free  metals,  suspended  in  colloidal  form  in  the  glass,  and  confer  a 
deep-red  color  upon  it.  Milk-glass  contains  finely  powdered  calcium 
phosphate  or  cryolite  in  suspension,  and  white  enamels  are  made  by 
adding  stannic  oxide. 

Glass  is  a  typical  amorphous  substance  (cf.  p.  154).  From  the 
facts  that  it  has  no  crystalline  structure,  and  that  it  softens  gradually 
when  warmed,  instead  of  showing  a  definite  melting-point,  it  is  re- 
garded as  a  supercooled  liquid  of  extreme  viscosity.  Most  single 
silicates  crystallize  easily,  and  have  definite  freezing-  (and  melting-) 
points.  Glass  may  be  regarded  as  a  solution  of  several  silicates. 
When  kept  for  a  considerable  length  of  time  at  a  temperature  in- 
sufficient to  render  it  perfectly  fluid,  some  of  its  components  crystallize 
out,  the  glass  becomes  opaque,  and  " devitrification"  is  said  to  have 
occurred.  The  absence  of  such  changes  in  cold  glass  may  be  attrib- 
uted to  that  general  hampering  of  all  molecular  movements  and 
interactions  which  is  characteristic  of  low  temperatures.  The  word 
" crystal"  popularly  applied  to  glass  is  thus  definitely  misleading. 
See  quartz-glass  (p.  634). 

Calcium-ion  Ca"1"1":  Analytical  Reactions.  —  Ionic  calcium 
is  colorless.  It  is  bivalent,  and  combines  with  negative  ions.  Many 
of  the  resulting  salts  are  more  or  less  insoluble  in  water.  Upon  the 
insolubility  of  the  carbonate,  phosphate,  and  oxalate  are  based 
tests  for  calcium-ion  in  qualitative  analysis  (see  p.  732).  The  pres- 
ence of  the  element  is  most  easily  recognized  by  the  brick-red  color 
its  compounds  confer  on  the  Bunsen  flame,  and  by  two  bands  —  a 
red  and  a  green  one  —  which  are  shown  by  the  spectroscope 
(p.  676). 


728  INORGANIC   CHEMISTRY 

STRONTIUM  SB 

The  compounds  of  strontium  resemble  closely  those  of  calcium, 
both  in  physical  properties  and  in  chemical  behavior. 

Occurrence.  —  The  carbonate  of  strontium  SrC03  is  found  as 
strontianite  (Strontian,  a  village  in  Argyleshire),  and  is  isomorphous 
with  aragonite.  The  sulphate,  celestite  SrSO4,  is  more  plentiful.  It 
shows  rhombic  crystals  which  are  isomorphous  with  those  of  anhydrite, 
often  have  a  blue  color,  and  are  commonly  associated  with  native 
sulphur  in  specimens  from  Sicily.  The  metal  may  be  isolated  by 
electrolysis  of  the  molten  chloride. 

Compounds  of  Strontium.  —  The  compounds  are  all  made 
from  the  natural  carbonate  or  sulphate.  The  former  may  be  dissolved 
directly  in  acids,  and  the  latter  is  first  reduced  by  means  of  carbon  to 
the  sulphide,  and  then  treated  with  acids. 

Strontium  chloride  SrCl2,6H2O,  made  in  one  of  the  above  ways, 
is  deposited  from  solution  as  the  hexahydrate.  The  nitrate  Sr(N03)2 
comes  out  of  hot  solutions  in  octahedrons  which  are  anhydrous. 
From  cold  water  the  tetrahydrate  is  obtained  (see  under  Manganous 
sulphate).  The  anhydrous  nitrate  is  mixed  with  sulphur,  charcoal, 
and  potassium  chlorate  to  make  "red  fire."  The  oxide  SrO  may  be 
secured  by  igniting  the  carbonate,  but  on  account  of  the  low  dissocia- 
tion tension  of  the  compound  it  is  obtained  with  greater  difficulty 
than  is  calcium  oxide  from  calcium  carbonate.  It  is  made  by  heating 
the  nitrate  or  hydroxide. 

Strontium  hydroxide  Sr(OH)2  is  made  by  heating  the  carbonate  in 
a  current  of  superheated  steam:  SrCOs  +  H2O  — >  Sr(OH)2  +  C02. 
This  action  takes  place  more  easily  than  does  the  mere  dissociation 
of  the  carbonate,  because  the  formation  of  the  hydroxide  liberates 
energy,  and  this  partially  compensates  for  the  energy  which  has  to 
be  provided  to  decompose  the  carbonate  (cf.  p.  278).  The  lowering 
of  the  partial  pressure  of  the  carbon  dioxide  by  the  steam  also  con- 
tributes to  the  result  (cf.  p.  709). 

The  hydrate  crystallizes  from  water  as  Sr(OH)2,8H2O,  and  is 
employed  in  separating  crystallizable  sugar  from  molasses.  By 
evaporation  of  the  extract  from  the  sugar-cane  or  beet-root,  as  much 
of  the  sugar  as  possible  is  first  secured  by  crystallization.  Then  the 
molasses  which  remains  is  mixed  with  a  saturated  solution  of  stron- 
tium hydroxide.  The  resulting  precipitate  of  sucrate  of  strontium, 
,  or  Ci2H22On,2SrO,  is  separated  by  a  filter-press,  made 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS       729 

into  a  paste  with  water,  and  treated  with  carbon  dioxide.  A  second 
nitration  parts  the  insoluble  carbonate  of  strontium  from  the  solu- 
tion of  sugar,  and  the  latter  is  evaporated  and  allowed  to  crystallize. 
Calcium  hydroxide,  which  gives  a  tricalcium  sucrate,  is  often  em- 
ployed in  the  same  way. 

Strontium-ion  Sr++  is  bivalent,  and  gives  insoluble  compounds 
with  carbonate-ion,  sulphate-ion,  and  oxalate-ion.  The  presence  of 
strontium  is  recognized  by  the  carmine-red  color  which  its  compounds 
give  to  the  Bunsen  flame  (see  also  p.  732).  Its  spectrum  shows 
several  red  bands  and  a  very  characteristic  blue  line. 

BARIUM  BA 

The  physical  and  chemical  properties  of  the  compounds  of  barium 
recall  those  of  strontium  and  calcium.  All  the  compounds  of  barium 
which  are  soluble  in  water,  or  can  be  brought  into  solution  by  the 
weak  acids  of  the  digestive  fluids,  are  poisonous. 

Occurrence.  —  Like  strontium,  barium  is  found  in  the  form  of 
the  carbonate,  witherite  BaC03,  which  is  rhombic  and  isomorphous 
with  aragonite,  and  the  sulphate  BaSO4,  heavy-spar  or  barite  (Gk. 
fiapvs,  heavy),  which  is  rhombic  and  isomorphous  with  anhydrite. 
The  density  of  the  sulphate  is  4.5,  while  that  of  other  compounds  of 
the  light  metals  does  not  generally  exceed  2.5.  The  free  metal, 
which  is  silver-white,  may  be  obtained  by  electrolysis  of  the  molten 
chloride. 

The  compounds  are  made  by  treating  the  natural  carbonate  with 
acids  directly,  or  by  first  reducing  the  sulphate  with  carbon  to 
sulphide,  or  converting  the  carbonate  into  oxide,  and  then  treating 
the  products  with  acids. 

Barium  Carbonate  BaCO^. —  The  precipitated  form  of  the 
carbonate  is  made  by  adding  sodium  carbonate  to  the  aqueous  ex- 
tract from  crude  barium  sulphide  (q.v.).  The  compound  is  also 
obtained  by  fusing  pulverized  barite  with  excess  of  sodium  carbonate, 
and  dissolving  the  sodium  salts  out  of  the  residue. 

This  carbonate  demands  a  high  temperature  (about  1500°)  for 
the  attainment  of  a  sufficient  dissociation  tension,  and  is  apt  then 
to  be  partially  protected  from  decomposition  by  the  melting  of  the 
oxide.  It  is  therefore  heated  with  powdered  charcoal  (cf.  p.  577) : 


730  INORGANIC  CHEMISTRY 

The  Sulphate  and  Sulphide.  —  The  natural  sulphate  BaS04 
is  the  source  of  many  of  the  compounds  of  barium.  The  precipitated 
sulphate,  made  by  adding  sulphuric  acid  to  the  aqueous  extract 
from  barium  sulphide,  is  used  in  making  white  paint  ("  permanent 
white"),  in  filling  paper  for  glazed  cards,  and  sometimes  as  an  adul- 
terant of  white  lead.  A  mixture  of  barium  sulphate  and  zinc  sul- 
phide ZnS,  prepared  in  a  special  way,  is  called  lithopone: 

BaS  +  ZnS04-»BaS04i  + 


Made  into  paint,  it  has  greater  covering  power  than  white  lead, 
does  not  darken  with  hydrogen  sulphide  as  does  the  latter,  and  is 
non-poisonous.  The  salt  is  highly  insoluble  in  water  and  is  hardly 
at  all  affected  by  aqueous  solutions  of  chemical  agents.  It  is  some- 
what soluble  in  hot,  concentrated  sulphuric  acid,  and  the  solution 
yields  crystals  of  a  compound  BaSO4,H2S04,  or  Ba(HSO4)2.  Calcium 
and  strontium  sulphates  behave  in  the  same  way.  All  three  com- 
pounds are  decomposed  by  water,  and  give  the  insoluble  sulphates. 
Barium  sulphide  BaS,  like  the  sulphides  of  calcium  and  strontium 
(p.  421),  is  slightly  soluble  in  water,  but  slowly  passes  into  solution 
owing  to  hydrolysis  and  formation  of  the  hydroxide  and  hydro- 
sulphide.  It  is  made  by  heating  the  pulverized  sulphate  with 
charcoal  : 

BaSO4  +  4C  -»  BaS  +  4CO. 

The  Chloride  and  Chlorate.  —  The  chloride  BaCl2  is  generally 
manufactured  by  heating  the  sulphide  with  calcium  chloride.  The 
whole  treatment  of  the  heavy-spar  is  carried  out  in  one  operation: 

BaS04  +  4C  +  CaCl2  -*  4CO  +  BaCl2  +  CaS. 

By  rapid  extraction  with  water,  the  chloride  can  be  separated  from 
the  calcium  sulphide  before  much  decomposition  of  the  latter  (cf. 
p.  687)  has  taken  place.  Barium  chloride  crystallizes  in  rhombic 
tables  as  a  dihydrate  BaCl2,2H20.  Aside  from  the  difference  in  com- 
position, this  compound  differs  from  the  ordinary  hydrated  chlorides 
of  calcium  and  strontium  in  being  non-hygroscopic  and  in  being 
capable  of  dehydration  by  heat  without  the  formation  of  any  hydro- 
gen chloride  (cf.  p.  655). 

Barium  chlorate  Ba(C103)2  is  made  by  treating  the  precipitated 
barium  carbonate  with  a  solution  of  chloric  acid.  It  is  deposited 
in  beautiful  monoclinic  crystals,  and  is  used  with  sulphur  and  char- 
coal in  the  preparation  of  "  green  fire." 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      731 

The  Oxides  and  Hydroxide.  —  The  oxide  of  barium  BaO  is 
manufactured  from  the  carbonate  (see  above)  or  sulphide.  In  the 
latter  case,  moist  carbon  dioxide  is  passed  over  the  sulphide,  and  the 
resulting  carbonate  is  then  treated  with  steam.  It  may  be  obtained 
in  pure  form  by  heating  the  nitrate.  The  oxide  unites  vigorously 
with  water  to  form  the  hydroxide.  When  heated  in  a  stream  of  air 
or  oxygen  it  gives  the  dioxide  BaO2.  This  change  and  its  reversal 
constitute  the  basis  of  Brin's  process  for  obtaining  oxygen  from  the 
air  (p.  82).  To  protect  the  oxide  from  conversion  into  the  carbon- 
ate and  hydrate,  which  are  not  decomposable  at  the  temperature 
employed,  the  air  must  be  carefully  purified  from  carbon  dioxide 
and  moisture  (cf.  p.  685). 

Barium  peroxide  BaO2,  when  made  by  union  of  oxygen  with  the 
monoxide,  is  a  compact  gray  mass.  A  hydrated  form  is  thrown 
down  as  a  crystalline  precipitate  when  hydrogen  peroxide  solution  is 
added  to  a  solution  of  barium  hydroxide: 

Ba(OH)2  +  H202  *=?  Ba02J  +  2H20. 

The  crystals  have  the  formula  BaO2,8H2O.  Similar  hydrates  of  the 
peroxides  of  strontium  and  calcium  may  be  made  in  the  same  way. 
In  all  three  cases  the  pure  peroxides  are  obtained  as  white  powders  by 
removal  of  the  water  of  hydration  by  very  gentle  heating  in  vacuo 
(cf.  p.  484) .  The  peroxides  of  strontium  and  calcium  are  not  formed 
by  direct  union  of  oxygen  with  the  oxides.  Barium  peroxide  is  used 
in  the  manufacture  of  hydrogen  peroxide  (p.  317). 

Barium  hydroxide  Ba(OH)2,  is  the  most  soluble  of  the  hydroxides 
of  this  group,  and  gives,  therefore,  the  highest  concentration  of 
hydroxide-ion.  The  solution  is  known  as  " baryta-water."  It 
is  also  the  most  stable  of  the  three  hydroxides,  and  may  be  melted 
without  decomposition.  A  hydrate  Ba(OH)2,8H2O  crystallizes 
on  cooling  a  saturated  solution.  It  is  much  used  in  quantitative 
analysis  for  making  standard  alkali-solutions.  Solutions  of  sodium 
or  potassium  hydroxide  may  acquire  varying  proportions  of  carbonate 
by  the  action  of  carbon  dioxide  from  the  air,  and  their  action  on 
indicators  loses  thereby  in  sharpness.  With  barium  hydroxide  this 
is  impossible,  for  the  carbonate  is  insoluble,  and  is  precipitated  from 
the  solution. 

Barium  Nitrate  Ba(]YO3)2.  —  This  salt  is  made  by  the  action  of 
nitric  acid  on  the  sulphide,  oxide,  hydroxide,  or  carbonate  of  barium. 
It  crystallizes  from  aqueous  solution  without  water  of  hydration. 


732  INORGANIC  CHEMISTRY 

Analytical  Reactions  of  the  Calcium  Family.  —  Barium-ion 
Ba++  is  a  colorless,  bivalent  ion.  Many  of  its  compounds  are  in- 
soluble in  water,  and  the  sulphate  is  insoluble  in  acids  also.  The 
spectrum  given  by  the  salts  contains  a  number  of  green  and  orange 
lines. 

In  solutions  of  salts  of  calcium,  strontium,  and  barium,  the  ions 
may  be  distinguished  by  the  fact  that  calcium  sulphate  solution  will 
precipitate  the  strontium  and  barium  as  sulphates,  but  will  leave 
salts  of  calcium  unaffected.  Similarly,  strontium  sulphate  solution 
precipitates  barium  sulphate,  and  does  not  give  any  result  with  salts 
of  the  two  first.  The  chromate  of  barium  is  precipitated  in  pres- 
ence of  acetic  acid,  while  the  chromates  of  strontium  and  calcium 
are  not  (cf.  p.  714),  and  there  are  other  differences  of  a  like  nature 
in  the  solubilities  of  the  salts. 

Exercises.  —  1.  Arrange  the  chromates  of  the  metals  of  this 
family  in  the  order  of  solubility  (see  Table).  Compare  the  solu- 
bilities with  those  of  the  carbonates,  oxalates,  and  sulphates  of  the 
metals  of  the  same  family. 

2.  What  must  be  the  approximate  total  molar  concentration  of 
the  solution  of  calcium  chloride  freezing  at  —48°  (p.  335)? 

3.  What  is  meant  by  fluorescence  (cf.  any  book  on  physics)? 

4.  What  will  be  the  ratio  by  volume,  at  150°,  of  the  nitrogen  per- 
oxide and  oxygen  given  off  by  the  decomposition  of  calcium  nitrate? 
What  would  be  the  nature  of  the  difference  between  the  ratio  at  150° 
and  that  at  room  temperature? 

5.  What  fact  about  the  heat  of  solution  of  gypsum  is  indicated  by 
its  change  of  solubility  with  temperature  (p.  717)? 

6.  What  is  the  significance  of  the  fact  that  hydrated  barium 
chloride  gives  no  hydrogen  chloride  when  heated? 

7.  What  are  the  advantages  of  removing  water  of  hydration  in 
vacuo  (p.  731)? 

8.  Explain  in  terms  of  the  ionic  hypothesis  the  precipitation  of 
the  sulphate  of  strontium  by  calcium  sulphate  solution,  and  the  ab- 
sence of  precipitation  when  the  latter  is  added  to  the  solution  of  a 
soluble  salt  of  calcium. 

9.  Construct  a  table  for  the  purpose  of  comparing  the  properties 
of  the  free  elements  of  this  family  and  also  the  properties  of  their 
corresponding  compounds. 

10.  Are  the  elements  of  this  family  typical  metals?     If  not,  in 
what  respects  do  they  fall  short  (p.  645)? 


METALLIC  ELEMENTS  OF  THE  ALKALINE  EARTHS      733 

11.  Apply  the  rule  of  precipitation  to  the  case  of  adding  sodium 
carbonate  to  a  solution  of  barium  chloride. 

12.  What  inference  do  you  draw  from  the  fact  that  the  chro- 
mates  of  calcium  and  strontium  are  not  precipitated  in  presence  of 
acetic  acid,  while  the  chromate  of  barium  is  so  precipitated?     Is 
the  inference  confirmed  by  reference  to  the  solubility  data? 

13.  Explain  the  fact  that  strontium  and  calcium  chromates  are 
easily  dissolved  by  acetic  acid,  while  barium  chromate  is  dissolved 
only  by  active  mineral  acids. 

14.  Explain  the  fact  that    all  the  carbonates,  save  those    of 
sodium,  potassium,  and  thallium,  are  precipitated  in  neutral  solu- 
tions,  but  not  in   acidified   solutions.     Why  is  the  precipitation 
incomplete  when  carbon  dioxide  is  led  through  solutions  of  salts 
of  the  metals,  but  more  complete  when  the  hydroxides  of  the  metals 
are  used? 


CHAPTER  XXXVI 
COPPER,    SILVER,    GOLD 

THE  three  metals  of  this  family,  being  found  free  in  nature,  are 
amongst  those  which  were  known  in  early  times.  They  are  the 
metals  universally  used  for  coinage  and  for  ornamental  purposes. 
They  are  the  three  best  conductors  of  electricity  (p.  645),  and  each 
represents  the  maximum  of  conductivity  in  the  periodic  series  to 
which  it  belongs.  In  malleability  and  ductility  silver  is  inter- 
mediate between  gold  and  copper  (p.  643),  but  in  electrical  conduc- 
tivity it  excels  both. 

The  Chemical  Relations  of  the  Copper  Family.  —  Copper 
(Cu,  at.  wt.  63.57),  silver  (Ag,  at.  wt.  107.88),  and  gold  (Au,  at.  wt. 
197.2),  occupy  the  right  side  in  the  second  column  of  the  table  of  the 
periodic  system  (opposite  inside  of  rear  cover),  and  the  chemical 
relations  (p.  226)  of  these  elements  are  in  many  ways  in  sharp  con- 
trast to  those  of  the  alkali  metals,  their  neighbors,  on  the  left  side: 

ALKALI  METALS  COPPER,  SILVER,  GOLD 

Very  active;  rapidly  oxidized  by  air;  Amongst  least  active  metals;  only 
displace  all  other  metals  from  com-  copper  is  oxidized  by  air;  displaced 
bination  (E.  M.  series,  p.  404).  by  most  other  metals.  Hence,  found 

free  in  nature  (p.  404). 

All  univalent  and  give  but  one  series     Cu1   and   Cu  :  two  series.     Ag1  :  one 

of  compounds.     Halides  all  soluble         series.     Au1  and  Au111  :  two  series. 

in  water.  Halides  of  univalent  series  insoluble. 

Oxides  and  hydroxides  strongly  basic,      Oxides    and    hydroxides    feebly    basic 

and  halides  not  hydrolyzed  (p.  648).         (except   Ag2O);    halides  hydrolyzed 

(except  Ag-halides).      Hence,   basic 
salts  are  numerous. 

Never  found  in  anion.  Give  no  com-  Frequently  in  anion,  e.g.,  K.Cu(CN)2, 
plex  cations.  K.Ag(CN)2,  K.AuO2,  K.Au(CN)2, 

and      in      complex      cation,      e.g., 
Ag(NH3)2.OH  and  Cu(NH3)4.(OH)2. 

On  account  of  their  inactivity  towards  oxygen,  and  their  easy 
recovery  from  combination  by  means  of  heat,  silver  and  gold,  to- 
gether with  the  platinum  family,  are  known  as  the  "noble  metals." 

734 


COPPER,   SILVER,   GOLD  735 

Univalent  copper  and  gold  resemble  in  some  ways  Hg1  and  Tl1, 
while  bivalent  copper  resembles  Zn11,  Mn11,  Fe11,  and  Nin,  and 
trivalent  gold  resembles  Ptiv.  This  family  is,  in  fact,  not  homoge- 
neous, and  the  close  relation  which,  amongst  metals,  subsists  between 
valence  and  chemical  properties  makes  comparisons  with  elements  of 
entirely  different  families  often  the  most  suggestive. 

COPPER  Cu 

Chemical  Relations  of  the  Element.  —  Copper  is  the  first 
metallic  element  showing  two  valences  which  we  have  encountered. 
In  such  cases  two  more  or  less  complete,  independent  series  of  salts 
are  known.  These  are  here  distinguished  as  cuprous  (univalent) 
and  cupric  (bivalent)  salts.  The  methods  by  which  a  compound 
of  one  series  may  be  converted  into  the  corresponding  compound  of 
the  other  series  should  be  noted. 

The  chief  cuprous  compounds  are  Cu2O,  CuCl,  CuBr,  Cul,  CuCN, 
Cu2S.  The  cuprous  compound  is  in  each  case  more  stable  (p.  148) 
than  the  corresponding  cupric  compound,  and  is  formed  from  it 
either  by  spontaneous  decomposition,  as  in  the  cases  of  the  iodide 
and  cyanide  (2CuI2  — •» 2CuI  +  12),  or  on  heating.  The  cuprous 
halides  and  cyanide  are  colorless  and  insoluble  in  water.  Cuprous- 
ion  Cu+  seems  to  be  colorless.  The  cuprous  salts  of  the  oxygen 
acids  are  unstable. 

The  familiar  cupric  compounds  are  more  numerous,  as  they  in- 
clude also  stable  salts  of  oxygen  acids,  like  CuSO4,  Cu(NOs)2,  etc. 
CuI2  and  Cu(CN)2  cannot  be  obtained  in  pure  form,  as  they  decom- 
pose, giving  the  cuprous  salts.  The  anhydrous  salts  are  usually 
colorless  or  yellow,  but  cupric  ion  Cu++  is  blue,  and  so,  therefore, 
are  the  aqueous  solutions  of  the  salts.  The  cupric  are  more  familiar 
than  the  cuprous  compounds,  since  cupric  oxide,  sulphate,  and 
acetate  are  the  compounds  of  copper  which  most  frequently  find 
employment  in  chemistry  and  in  the  arts.  All  the  soluble  salts  of 
copper  are  poisonous. 

In  electrolyzing  salts  of  copper,  a  given  amount  of  electricity  will 
deposit  twice  as  much  copper  from  a  cuprous  salt  as  from  a  cupric 
salt  (p.  350),  since  cuprous-ion  carries  only  half  as  great  a  charge, 
weight  for  weight,  as  cupric-ion. 

Writers  on  chemistry  still  (p.  137)  frequently  double  (Cu2Cl2, 
etc.)  the  formulae  of  cuprous  salts.  The  molecular  weights  in  organic 
solvents  (cf.  p.  335),  however,  in  many  cases  accord  with  the  simple 


736  INORGANIC  CHEMISTRY 

formulae.  Some  higher  molecular  weights  observed  in  solution  and 
the  vapor  density  of  cuprous  chloride  (6.6,  corresponding  nearly  to 
Cu2Cl2)  might  be  regarded  as  being  due  to  association  (imperfect 
dissociation,  cf.  p.  282).  The  formation  of  numerous  double  or 
complex  compounds  like  HCl,CuCl  (  =HCuCl2),  which  may  be 
regarded  as  acid  salts,  however,  lends  support  to  the  view  that  the 
formulae  should  be  doubled.  Inasmuch  as  the  behavior  of  the 
salts  is  sufficiently  well  represented  by  the  simple  formulae,  these 
are  here  used  throughout  (see  under  Silver,  p.  749). 

In  addition  to  (1)  having  two  valences  Cu1  and  Cu11,  and  there- 
fore two  series  of  compounds  (two  oxides,  two  chlorides,  etc.),  each 
of  these  states  of  copper  also  joins  with  other  elements  to  form  (2) 
complex  positive  ions  such  as  Cu(NH3)2+  and  Cu(NH3)4++,  just  as 
hydrogen  and  nitrogen  form  the  complex  positive  ion  NH4+,  and 
the  univalent  form  also  gives  (3)  stable  complex  negative  ions  such 
as  Cu(CN)2~,  CuCl2~.  None  of  the  metallic  elements  discussed  in 
the  two  preceding  chapters  showed  any  of  these  peculiarities.  Many 
of  the  metals  to  be  discussed  later  exhibit  one  or  more  of  them, 
however.  Especial  attention  should  therefore  be  given  to  the  chem- 
istry of  copper,  in  order  that  the  behavior  which  such  relations 
entail  may  be  mastered  at  the  first  encounter,'  and  the  same  rela- 
tions may  be  instantly  recognized  and  understood  when  they  re- 
appear in  other  connections. 

There  is  only  one  other  peculiarity  which  a  metallic  element 
frequently  shows,  although  copper  does  not  exhibit  it.  This  is  (4) 
the  ability  of  its  hydroxide  to  be,  not  only  basic,  as  metallic  hydrox- 
ides by  definition  (p.  646)  must  be,  but  also  acidic.  This  be- 
havior we  encounter  first  in  the  case  of  gold  (see  p.  760)  and  in 
simpler  and  more  familiar  form  in  the  case  of  zinc  (see  next 
chapter). 

Occurrence.  —  Copper  is  found  free  in  the  Lake  Superior  region, 
in  China,  and  in  Japan.  The  sulphides,  copper  pyrites  CuFeS2  and 
chalcocite  Cu2S,  are  worked  in  Montana,  in  southwest  England,  and 
in  Spain.  Malachite,  Cu2(OH)2CO3  (=  Cu(OH)2,CuCO3),  and  az- 
urite,  Cu3(OH)2(CO3)2  (  =  Cu(OH)2,2CuC03),  both  basic  carbon- 
ates, are  mined  in  Arizona,  Siberia,  and  elsewhere.  Cuprite  or 
ruby  copper  Cu20  is  also  an  important  ore.  The  name  of  the  element 
comes  from  the  fact  that  in  ancient  times  copper  mines,  long  since 
worked  out,  existed  in  Cyprus.  The  element  is  found  in  the  feathers 
of  some  birds,  in  the  haemocyanin  of  the  blood  of  the  cuttle-fish, 


COPPER,   SILVER,  GOLD  737 

which  is  blue  when  arterial  and  colorless  when  venous,  and  elsewhere 
in  living  organisms. 

Extraction  from  Ores.  —  For  isolating  native  copper  it  is  only 
necessary  to  separate  the  metal,  by  grinding  and  washing,  from  the 
rock  through  which  it  ramifies,  and  to  melt  the  almost  pure  powder  of 
copper  with  a  flux  (p.  652).  The  carbonate  and  oxide  ores  require 
coal,  in  addition,  for  the  removal  of  the  oxygen. 

The  liberation  of  copper  from  the  sulphide  ores  is  difficult,  and 
often  involves  very  elaborate  schemes  of  treatment.  This  arises 
from  the  fact  that  many  copper  ores  contain  a  large  amount  of  the 
sulphides  of  iron,  and  these  have  to  be  removed  by  conversion  into 
oxide  (by  roasting)  and  then  into  silicate  (with  sand).  The  silicate 
forms  a  flux,  and  separates  itself  from  the  molten  copper  sulphide 
(" matte").  In  Montana  it  is  found  possible  to  abbreviate  the  treat- 
ment. The  ore  is  first  roasted  until  partially  oxidized.  It  is  then 
melted  in  a  cupola  or  a  reverberating  furnace,  and  placed  in  large  iron 
vessels  like  Bessemer  converters  (q.v.)  provided  with  a  lining  rich  in 
silica.  A  blast  of  air  mixed  with  sand  is  now  blown  through  the 
mass.  The  iron  is  completely  oxidized  to  FeO  and  made  into  silicate 
FeSiOs,  the  sulphur  escapes  as  sulphur  dioxide,  and  arsenic  and  lead 
are  likewise  removed  by  this  treatment.  The  silicate  of  iron  floats 
as  a  slag  upon  the  copper  when  the  contents  of  the  converter  are 
poured  out.  The  resulting  copper  is  pure  enough  to  be  cast  in  large 
plates  and  purified  by  electrolysis  (see  p.  747). 

Much  copper  ore  is  of  low  grade,  containing  perhaps  only  2  per 
cent  of  copper  ore  and  98  per  cent  of  rock  material.  From  such 
ores  the  usual  methods  of  washing  often  recover  only  70  per  cent 
or  less  of  the  copper  ore  present,  and  30  per  cent  or  more  is  lost. 
The  froth  flotation  process  raises  the  proportion  recovered  to  85  or 
90  per  cent  of  the  whole.  The  finely  crushed  ore  is  agitated  with 
water,  to  which  is  added  some  cheap  oil  and  sometimes  a  little 
sulphuric  acid.  The  mixture  is  then  allowed  to  flow  into  a  larger 
tank  of  water,  in  which  the  rock  material  immediately  sinks  to  the 
bottom  while  the  particles  of  ore  are  contained  in  the  oily  froth 
which  rises  to  the  top.  The  plant  also  occupies  less  than  one-tenth 
of  the  space,  and  uses  less  than  half  the  power  required  for  treating 
the  same  amount  of  ore  by  washing. 

The  world's  production  (1913)  is  about  a  million  metric  tons,  of 
which  the  United  States  furnished  58  per  cent,  South  America  11, 
Japan  6,  and  Germany  4.  The  proportions  of  the  whole  consumed 


738  INORGANIC  CHEMISTRY 

were,  approximately,  United  States  33  per  cent,  Great  Britain  22 
per  cent,  Germany  22  per  cent. 

Physical  Properties.  —  Copper  is  red  by  reflected  and  greenish 
by  transmitted  light.  Native  copper  shows  crystals  of  the  regular 
system  (p.  172).  It  melts  at  1083°,  and  therefore  much  more  easily 
than  pure  iron  (1530°).  Its  density  is  8.93.  When  steel  draw- 
plates  are  used,  it  can  be  drawn  into  wire  with  a  diameter  of  only 
0.2  mm.,  and  by  means  of  plates  provided  with  perforated  diamonds 
the  diameter  of  the  wire  can  be  reduced  to  0.03  mm.  (1  kilometer 
weighs  only  7  g.).  The  metal  after  drawing  is  more  tenacious,  but 
conducts  electricity  less  well.  Copper  has  a  transition  point  at  71.7°, 
and  shows  different  physical  properties  above  that  point. 

Chemical  Properties.  —  In  dry  oxygen,  copper  does  not  rust. 
In  moist  oxygen  a  thin  film  of  cuprous  oxide  is  formed,  and  in  ordinary 
air  a  green  basic  carbonate  (not  verdigris,  q.v.).  It  does  not  decom- 
pose water  at  any  temperature  or  displace  hydrogen  from  dilute  acids 
(p.  404).  On  the  other  hand,  hydrogen,  absorbed  in  platinum  or 
even  in  charcoal,  liberates  copper:  Cu++  +  H2  — >  Cu  -f  2H+,  when 
immersed  in  solutions  of  copper  salts.  The  metal  attacks  oxygen 
acids  (pp.  425,  535),  however.  Again,  acids  like  hydrochloric  acid, 
in  conjunction  with  oxygen  from  the  air,  do  act  slowly  upon  copper: 
2Cu  +  4HC1  +  O2  — >  2CuCl2  +  2H2O.  This  sort  of  simultaneous 
action  of  two  agents  is  frequently  used,  as  in  making  silicon  tetra- 
chloride  (p.  632).  In  a  similar  way  sea-water  and  air  slowly  corrode 
the  copper  sheathings  of  ships,  giving  a  basic  chloride,  Cu4(OH)6Cl2- 
H2O  ( =  3Cu(OH)2,CuCl2,H2O),  which  is  found  in  nature  as  atakamite. 

On  account  of  its  resistance  to  the  action  of  acids,  copper  is  used 
for  many  kinds  of  vessels,  for  covering  roofs  and  ships'  bottoms,  and 
for  coins.  It  furnishes  also  electrotype  reproductions  of  medals,  of 
engraved  plates,  of  type,  etc.  (see  p.  746).  In  the  mines  near  Butte, 
the  presence  of  sulphuric  acid  in  the  water,  which  attacks  steel, 
compels  the  use  of  pumping  machinery  of  copper. 

Alloys.  —  The  qualities  of  copper  are  modified  for  special  pur- 
poses by  alloying  it  with  other  metals.  Brass  contains  18-40  per 
cent  of  zinc,  and  melts  at  a  lower  temperature  (p.  644)  than  does 
copper.  A  variety  with  little  zinc  is  beaten  into  thin  sheets,  giving 
Dutch-metal  (" gold-leaf ").  Bronze  contains  3-8  per  cent  of  tin,  11 
or  more  per  cent  of  zinc,  and  some  lead/  It  was  used  for  making 


COPPER,  SILVER,  GOLD  739 

weapons  and  tools  before  means  of  hardening  iron  were  known,  and 
later,  on  account  of  its  fusibility,  continued  to  be  employed  for  cast- 
ings until  displaced  largely  by  cast-iron  (discovered  in  the  eighteenth 
century).  For  works  of  art  it  is  preferred  to  copper  because  of  its 
fusibility,  its  color,  and  its  more  rapid  acquirement  of  a  much  prized 
" patina,"  due  to  surface  corrosion.  Artificial  " bronzing"  of  brass 
is  affected  by  applying  a  solution  of  arsenious  oxide  in  hydrochloric 
acid  (AsCla).  The  zinc  displaces  some  arsenic,  which  combines  with 
the  copper.  Brass  instruments  are  "bronzed"  by  means  of  a  dilute 
solution  of  chloroplatinic  acid  (q.v.),  from  which  the  zinc  displaces 
platinum.  Gun-metal  contains  10  per  cent,  and  bell-metal  20-24 
per  cent  of  tin.  German  silver  contains  19-44  per  cent  of  zinc  and 
6-22  per  cent  of  nickel,  and  shows  none  of  the  color  of  copper.  Alu- 
minium-bronze contains  5-10  per  cent  of  aluminium,  and  resembles 
gold  in  color.  When  it  contains  some  iron,  it  can  be  worked  at  a 
red  heat,  but  not  welded.  Silicon-bronze  contains  not  more  than  5 
per  cent  of  silicon,  and  is  made  by  adding  silicide  of  copper  (made  in 
the  electric  furnace,  p.  569)  to  copper.  It  has  usually  only  60 
per  cent  of  the  conductivity  of  pure  copper,  but  is  nearly  twice  as 
tenacious,  and  is  used  for  telephone  and  over-head  electric  wires. 
Phosphor-bronze  contains  copper  and  tin  (100  : 9)  with  J-l  part  of 
phosphorus,  and  is  employed  for  certain  parts  of  machines.  Ships' 
propellers  are  made  of  manganese-bronze  (30  per  cent  manganese). 
In  many  of  these  alloys  the  metals  are  partly  in  the  form  of  chemical 
compounds,  such  as  Cu3Sn  and  Cu2Zn3. 

Cupric  Chloride  CuClz.  —  This  compound  is  made  by  union 
of  copper  and  chlorine,  by  treating  the  hydrate  or  carbonate  with 
hydrochloric  acid,  or  by  heating  copper  with  hydrochloric  acid  and 
some  nitric  acid,  the  latter  being  used  simply  as  an  oxidizing  agent: 
Cu  +  2HC1  +  O  ->  CuCl2  +  H2O.  The  blue  crystals  of  a  hydrate, 
CuCl2,2H2O,  are  deposited  by  the  solution.  The  anhydrous  salt  is  yel- 
low. Dilute  solutions  are  blue,  the  color  o :  cupric-ion,  but  concentrated 
solutions  are  green  on  account  of  the  presence  of  the  yellow  molecules 
(p.  378).  The  aqueous  solution  is  acid  in  reaction  (p.  399).  When 
excess  of  ammonium  hydroxide  is  added  to  the  solution,  the  basic 
chloride,  cupric  oxychloride  Cu4(OH)6Cl2  (p.  738),  which  is  at  first 
precipitated,  redissolves,  and  a  deep-blue  solution  is  obtained. 
This  on  evaporation  yields  deep-blue  crystals  of  hydrated  ammonio- 
cupric  chloride  Cu(NH3)4.Cl2,H2O.  The  deep-blue  color  of  the 
solution,  which  is  given  by  all  cupric  salts,  is  that  of  ammonio- 


740  INORGANIC  CHEMISTRY 

cupric-ion  Cu(NH3)4++.  The  dry  salt  also  absorbs  ammonia,  giv- 
ing CuCl2,6NH3.  This  and  the  preceding  compound  have  an  appre- 
ciable tension  of  ammonia,  and  when  warmed,  or  placed  under 
reduced  pressure,  leave  first  CuCl2,2NH3,  and  finally  CuCl2. 

Cuprous  Chloride  CuCL  —  This  salt  is  formed  when  dry  cupric 
chloride  is  heated :  2CuCl2  ^  2CuCl  +  C12.  It  may  be  made  by 
boiling  cupric  chloride  solution  with  hydrochloric  acid  and  copper 
turnings : 

CuCl2  +  Cu  ->  2CuCl     or     Cu++  +  Cu  ->  2Cu+ 

The  solution  contains  compounds  of  cuprous  chloride  with  hydrogen 
chloride  HCl,CuCl  or  HCuCl2  and  H2CuCl3,  which  are  decomposed 
when  water  is  added.  The  cuprous  chloride  is  insoluble  in  water, 
and  forms  a  white  crystalline  precipitate. 

Cuprous  chloride  is  hydrolyzed  quickly  by  hot  water,  giving, 
finally,  red,  hydrated  cuprous  oxide,  Cu2O.  When  dry  it  is  not 
affected  by  light,  but  in  the  moist  state  becomes  violet  and,  finally, 
nearly  black.  The  action  is  said  to  be  2CuCl  — >  CuCl2  +  Cu,  the 
copper  being  adsorbed  as  a  colloid  (hence  the  color  changes) .  In  moist 
air  it  turns  green,  and  is  oxidized  to  cupric  oxychloride  (p.  739).  It 
is  dissolved  by  hydrochloric  acid,  giving  the  colorless  complex  acids 
HCuCl2  and  H2CuCl3.  The  solution  is  oxidized  by  the  air,  turning 
first  brown  and  then  green.  Cuprous  chloride  also  dissolves  in 
ammonium  hydroxide,  giving  ammonio-cuprous  chloride,  probably 
Cu(NH3)2.Cl,  the  ion  Cu(NH3)2+  being  colorless.  The  solution  is 
quickly  oxidized  by  the  air,  turns  deep-blue,  and  then  contains 
Cu(NH3)4++.  The  solution  of  cuprous  chloride  in  hydrochloric  acid  is 
used  for  absorbing  carbon  monoxide  from  gaseous  mixtures.  A 
crystalline  compound  (CuCO.Cl,2H20  ?)  has  been  isolated  from  the 
solution. 

The  Bromides  and  Iodide  of  Copper.  —  By  treatment  of  cop- 
per with  bromine-water,  and  slow  evaporation  of  the  solution, 
jet-black  crystals  of  anhydrous  cupric  bromide  CuBr2  are  obtained. 
A  concentrated  aqueous  solution  is  deep-brown  in  color,  and  the  grad- 
ual ioniza  tion  of  the  molecules  as  the  solution  is  diluted  is  well  shown 
by  this  salt  (p.  378).  The  ionization  is  here  accompanied  by  evolu- 
tion of  heat  (p.  368),  as  it  is  also  in  the  cases  of  cupric  chloride  and 
cupric  sulphate,  and  in  the  ionized  condition  the  substances  contain 
available  energy  than  in  the  molecular.  In  these  cases,  there- 


COPPER,  SILVER,  GOLD  741 

fore,  when  the  temperature  is  raised  the  ionization  diminishes 
(p.  305). 

When  cupric  bromide  is  heated,  bromine  is  given  off,  and  cuprous 
bromide  CuBr  remains. 

Cupric  iodide  CuI2  appears  to  be  unstable  at  ordinary  tempera- 
tures. When  a  soluble  iodide  is  added  to  a  cupric  salt,  a  white 
precipitate  of  cuprous  iodide  Cul  and  free  iodine  are  obtained: 


The  iodine  may  be  dissolved  in  excess  of  a  soluble  iodide  (p.  665),  or 
reduced  to  hydrogen  iodide  with  sulphurous  acid  (p.  445). 

The  Solution  of  Insoluble  Salts  when  Complex  Ions  are 
Formed.  —  The  solution  of  an  insoluble  salt  like  cuprous  chloride 
by  hydrochloric  acid  or  ammonium  hydroxide  is  typical  of  a  great 
variety  of  actions  of  which  we  here  meet  one  of  the  first  examples 
(cf.  p.  405).  The  explanation  involves  only  principles  already  used 
in  other  cases. 

Since  a  salt  is  normally  less  soluble  in  an  acid  having  the  same 
anion  (p.  699),  the  dissolving  of  cuprous  chloride  in  hydrochloric  acid 
requires  a  special  explanation,  namely,  the  fact  that  here  a  soluble  com- 
plex acid,  H.CuCl2  is  formed.  The  chloride-ion  of  the  hydrogen 
chloride  must  indeed  tend  to  repress  the  ionization  of  the  dissolved 
part  of  the  cuprous  chloride,  so  that  a  smaller  concentration  of 
Cu+  remains.  But  the  complex  negative  ion  CuCl2~  which  is  formed, 
is  very  little  dissociated,  and  gives  a  still  smaller  concentration  of 
Cu+  (CuCl2~  ^?  Cu+  +  2C1~).  The  ion-product  of  cuprous  chloride, 
and  the  concentration  relations  of  the  ionic  substance  CuCl2~~  and 
its  dissociation  products  (Cu+  and  2C1~)  are  symbolized  as  follows: 


The  value  of  [Cu+]  from  cuprous  chloride  (first  formula)  is,  in 
general,  greater  than  its  value  from  the  ion  CuCl2~  of  HCuCl2  (sec- 
ond formula),  when  excess  of  HC1  is  present.  Hence,  the  Cu+ 
tends  to  pass  over  into  the  more  stable  compound,  where  it  is  more 
completely  combined.  More  CuCl  dissolves  to  replace  the  Cu+ 
which  has  been  removed,  and  the  change  stops  when  the  CuCl  is 
all  dissolved,  or  the  values  of  [Cu+]  from  both  compounds  have 
become  equal.  Thus,  the  complex  ion  is  formed  at  the  expense  cf 


742  INORGANIC  CHEMISTRY 

the  Cu+  of  the  insoluble  cuprous  chloride,  and  the  latter  goes  into 
solution  progressively  in  the  effort  to  restore  the  balance: 

Cud  (solid)  t=r  CuCl  (dslvd)  fc*  Cl-   +  CU+  j  _  c  r,  _  (,  ,    ,s 
2HC1  *  -~ 


The  same  exact  laws  of  equilibrium  used  in  discussing  the  dissolving 
of  salts  by  acids  (p.  713)  may  be  applied  to  the  whole  procedure. 

Similar  behavior  is  shown  by  the  cyanides  of  copper,  silver,  iron, 
etc.  (q.v.),  of  which  many  complex  compounds  are  known. 

The  dissolving  of  cuprous  chloride  by  the  free  ammonia  of  ammo- 
nium hydroxide  is  explained  in  the  same  way.  The  only  difference 
is  that  here  the  copper  is  in  the  complex  positive  ion.  The  ion 
Cu(NH3)2+  gives  little  Cu+  —  less  than  does  cuprous  chloride,  in 
spite  of  the  insolubility  of  the  latter.  Hence  the  salt  passes  into 
solution  until  the  ion-product  [Cu+]  X  [Cl~]  with  continually 
increasing  [Cl~]  reaches  the  value  for  a  saturated  solution,  or  until 
the  solid  is  exhausted. 

The  deep-blue  colored  ion  Cu(NH3)4++  given  by  cupric  chloride 
and  other  cupric  salts  is  also  very  little  ionized.  Hence  ammonium 
hydroxide  interacts  with  all  the  insoluble  cupric  compounds  save 
only  cupric  sulphide,  which  is  the  most  insoluble  of  all  —  that  is, 
the  one  giving  the  smallest  concentration  of  cupric-ion.  Conversely, 
the  sulphide  is  the  only  insoluble  compound  of  copper  which  can  be 
precipitated  from  ammoniacal  solution.  Cupric  ferrocyanide  Cu2Fe- 
(CN)e,  however,  requires  a  large  excess  of  ammonium  hydroxide  for 
complete  interaction.  Zinc  and  other  more  active  metals,  however, 
slowly  precipitate  metallic  copper,  thereby  showing  that  some  cupric- 
ion  is  present. 

Foregoing  Explanation  Restated.  —  We  may  restate  the  ex- 
planation by  answering  a  question:  Why  does  cuprous  chloride 
interact  with,  and  go  into  solution  in  hydrochloric  acid?  Because 
it  forms  a  complex  compound  HCuCl2,  and,  with  the  concentrations 
usually  employed,  the  molecular  concentration  of  cuprous-ion  in  the 
solubility  product  of  cuprous  chloride  is  greater  than  the  molecular 
concentration  of  the  same  ion  in  the  solution  of  the  complex  com- 
pound. 

The  answer  in  other  cases  takes  the  same  form.  Thus,  for 
cupric  hydroxide  Cu(OH)2  dissolving  in  ammonium  hydroxide 
solution,  substitute  cupric  hydroxide  for  cuprous  chloride  and 
Cu(NH3)4(OH)2  for  HCuCl2. 


COPPER,  SILVER,   GOLD  743 

Cuprous  Oxide  Cu2O.  —  This  oxide  is  red_in  color,  and  natural 
specimens  show  octahedral  forms.  It  is  produced  by  oxidation  of 
finely  divided  copper  at  a  gentle  heat,  or  by  the  addition  of  bases  to 
cuprous  chloride,  and  is  best  made  by  the  action  of  glucose  (p.  605) 
on  cupric  hydroxide  (see  Fehling's  solution,  below).  The  latter 
is  reduced  by  the  former,  and  the  resulting  hydrated  cuprous  oxide 
forms  a  pale-brown  precipitate  which  quickly  becomes  bright  red. 
The  simple  hydroxide  CuOH  is  unknown,  but  the  above  mentioned 
precipitate  has  approximately  the  composition  4Cu2O,H20,  and  yields 
Cu2O  when  heated. 

Cuprous  oxide  is  acted  upon  by  hydrochloric  acid,  giving  cuprous 
chloride,  or  rather  HCuCl2.  It  also  dissolves  in  ammonium  hydrox- 
ide, giving,  probably,  Cu(NH3)2.OH,  which  is  colorless.  With 
dilute  oxygen  acids  part  of  it  is  oxidized,  giving  the  cupric  salt,  and 
part  is  reduced  to  metallic  copper: 

Cu2O  +  H2SO4  ->  CuS04  +  Cu  +  H20. 


Cupric  Oxide  and  Hydroxide.  —  Cupric  oxide  CuO  (black)  is 
formed  by  heating  copper  in  a  stream  of  oxygen,  or  by  igniting 
the  nitrate,  carbonate,  or  hydroxide.  Although  not  soluble  in  water, 
it  absorbs  moisture  from  the  air,  probably  because  it  is  porous  and 
has  much  surface.  When  heated  strongly  it  loses  some  oxygen, 
and  is  partly  reduced  to  cuprous  oxide.  Its  chief  use  is  in  the  analysis 
of  compounds  of  carbon.  When  heated  with  the  latter,  it  oxidizes 
the  hydrogen  to  water,  and  the  carbon  to  carbon  dioxide.  The 
operation  is  performed  in  a  tube  through  which  passes  a  stream  of 
oxygen,  and  the  products  are.  caught  in  glass  vessels  containing 
calcium  chloride  and  potassium  hydroxide,  respectively,  and  the 
increase  in  weight  of  each  is  determined. 

Cupric  hydroxide  Cu(OH)2  is  precipitated  as  a  gelatinous  sub- 
stance by  addition  of  sodium  or  potassium  hydroxide  to  a  solution 
of  a  cupric  salt:  Cu++  +  20H~  -»  Cu(OH)2.  When  the  mixture  is 
boiled,  the  hydroxide  loses  water  and  forms  a  black,  hydrated  cupric 
oxide  Cu(OH)2,2CuO  (?). 

The  hydroxide  interacts  with  ammonium  hydroxide,  forming 
the  soluble  compound  Cu(NH3)4.(OH)2,  which  imparts  a  deep-blue 
color  to  the  solution.  Various  forms  of  cellulose,  such  as  filter 
paper  and  cotton,  dissolve  in  this  solution  and  are  reprecipitated 
when  the  ammonium  hydroxide  is  neutralized  with  sulphuric  acid. 
Artificial  silk  is  made  by  pressing  the  solution  through  dies  into  the 
precipitant.  Paper  and  cotton  goods,  when  passed  first  through  one 


744  INORGANIC  CHEMISTRY 

and  then  the  other  of  these  liquids,  receives  a  tough,  waterproof 
surface. 

Cupric  hydroxide  interacts  with  a  solution  of  sodium  tartrate 
Na2.(C02)2C2H2(OH)2,  giving  a  deep-blue  liquid  (practically  "Fehling's 
solution").  In  this  action,  it  enters  into  the  negative  ion,  as  is  shown 
by  electrolysis,  interacting  apparently  with  the  hydroxyl  groups  of 
the  tartrate-ion.  The  solution  is  used  in  testing  for,  and  estimating 
quantities  of  glucose  (p.  605),  and  other  reducing  substances.  Cu- 
prous oxide  is  precipitated  (see  p.  743). 

Cupric  Nitrate  Cu(ArO3)2.  —  The  nitrate  is  made  by  treating 
cupric  oxide  or  copper  with  nitric  acid  (p.  535),  and  is  obtained  from 
the  solution  as  a  deliquescent,  crystalline  hydrate.  The  hexahydrate 
is  secured  at  temperatures  below  24.5°,  its  transition  point  (p.  689), 
and  the  trihydrate  from  24.5°  up  to  114.5°  (its  transition  point; 
see  under  Manganous  sulphate).  When  dehydrated  at  65°  the  salt 
is  partly  hydrolyzed,  and  a  basic  nitrate  Cu4(OH)6(N03)2  remains. 

Carbonate  of  Copper.  —  No  normal  carbonate  CuCO3  can  be 
obtained.  A  basic  carbonate  (malachite)  is  found  in  nature,  and  is 
precipitated  by  adding  soluble  carbonates  to  cupric  salts: 

2CuS04  +  2Na2C03  +  H20  ->  Cu2(OH)2C03  +  2Na2S04  +  C02. 


The  carbonate,  if  formed,  would  be  hydrolyzed  by  water  (p.  647). 

Cyanides  of  Copper.  —  With  potassium  cyanide  and  a  solu- 
tion of  a  cupric  salt,  cupric  cyanide  Cu(NC)2  is  precipitated.  This 
is  not  stable,  however,  and  gives  off  cyanogen,  leaving  cuprous 
cyanide: 

2Cu(CN)2  ->  2CuCN  +  C2N2T. 

Cuprous  cyanide  is  insoluble  in  water,  but  interacts  with  an  excess  of 
potassium  cyanide  solution,  producing  a  colorless  liquid,  from  which 
K.Cu(CN)2  (  =  KCN,CuCN)  potassium  cuprocyanide,  may  be  ob- 
tained in  colorless  crystals.  The  complex  anion  Cu(CN)2~  is  so 
little  ionized  to  Cu+  and  2CN~  that  all  insoluble  copper  compounds, 
including  cupric  sulphide,  interact  with  potassium  cyanide;  and 
none  of  them  can  be  precipitated  from  the  solution.  Zinc  is  actually 
unable  to  displace  copper  from  such  a  solution.  The  cause  of  the 
solution  of  the  salts  is  the  same  as  when  the  complex  ions  Cu(NH3)2+, 
Cu(NH3)4++  and  CuCl2-  are  formed  (p.  741). 


COPPER,   SILVER,   GOLD  745 

Cupric  Acetate.  —  By  the  oxidation  of  plates  of  copper,  sep- 
arated by  cloths  saturated  with  acetic  acid  (vinegar),  a  basic  acetate 
of  copper  (verdigris)  is  obtained: 

6Cu  +  8HC2H302  +  302  -»  2Cu3(OH)2(C2H302)4  +  2H20. 

It  is  used  in  manufacturing  green  paint,  is  insoluble  in  water,  and  is 
unaffected  by  light.  It  dissolves  in  acetic  acid,  and  green  crystals 
of  the  normal  acetate  Cu(C2H3O2)2,H20  are  obtained  from  the  solu- 
tion. The  basic  acetate  is  used  in  preparing  Paris  green.  A  hot 
solution  of  arsenious  acid  H3AsO3  is  mixed  with  a  paste  of  verdigris 
and  a  little  acetic  acid  and  boiled.  A  precipitate  of  Paris  green 
Cu(C2H3O2)2,Cu3(AsO3)2,  which  has  a  unique  light-green  color,  is  thrown 
down.  On  account  of  their  poisonous  nature,  this  compound  and 
Scheele's  green  CuHAsO3  are  little  used  as  pigments.  The  former 
is  chiefly  made  for  use  in  the  extermination  of  potato-beetles  and 
other  insects  and  for  employment  in  the  destruction  of  parasitic 
fungi. 

Cupric  Sulphate  CuSO±.  —  This  salt  is  obtained  by  heating 
copper  in  a  furnace  with  sulphur,  and  admitting  air  to  oxidize  the 
cuprous  sulphide.  The  mixture  of  cupric  sulphate  and  cupric 
oxide  which  is  formed  is  treated  with  sulphuric  acid.  The  salt  is 
also  made  by  allowing  dilute  sulphuric  acid  to  trickle  over  granulated 
copper  while  air  has  free  access  to  the  material :  2Cu  +  2H2SO4  +  02 
— >  2CuS04  H-  2H2O.  When  concentrated  and  at  a  high  temperature, 
sulphuric  acid  will  itself  act  as  the  oxidizing  agent  (cf.  p.  425). 

Cupric  sulphate  crystallizes  as  pentahydrate  CuSO4,5H2O  in  blue 
asymmetric  crystals  (Fig.  55,  p.  151),  and  in  this  form  is  called  blue- 
stone  or  blue  vitriol.  The  dissociation  of  this  hydrate  has  been 
discussed  on  page  153.  The  aqueous  solution  has  an  acid  reaction 
(p.  399).  The  anhydrous  salt  is  white,  and  can  be  crystallized  in  thin 
needles  (rhombic  system?)  from  solution  in  hot,  concentrated  sul- 
phuric acid  (cf.  pp.  151-154).  Cupric  sulphate  is  employed  in  copper- 
plating  (see  p.  747),  in  batteries,  and  as  a  mordant  in  dyeing  (q.v.). 
A  minute  proportion  is  added  to  drinking  water,  to  destroy  algcs, 
which  otherwise  propagate  in  the  reservoirs  and  give  a  disagreeable 
taste  and  odor  to  the  water.  The  seeds  of  cereals  are  moistened  with 
a  dilute  solution,  before  planting,  to  prevent  the  growth  of  fungi 
(smuts).  A  solution,  mixed  with  milk  of  lime  (Cu(OH)2  is  pre- 
cipitated), Bordeaux  mixture,  is  largely  used  as  a  spray  on  grape 
vines  and  other  plants  to  prevent  the  growth  of  fungi. 


746  INORGANIC  CHEMISTRY 

When  ammonium  hydroxide  is  added  to  cupric  sulphate  solution, 
a  pale-green  basic  sulphate  Cu4(OH)6SO4(?)  is  first  precipitated. 
With  excess  of  the  hydroxide,  the  blue  Cu(NH3)4++  ion  (p.  742)  is 
formed,  and  crystals  of  ammonio-cupric  sulphate  Cu(NH3)4.SO4,H20 
can  be  obtained  from  the  solution.  This  compound  easily  loses 
water  and  ammonia  (by  stages),  leaving  successively  CuSO4,2NH3 
and  CuSO4,NH3.  Cupric  sulphate  also  combines  with  potassium 
and  ammonium  sulphates,  giving  double  salts  of  the  form  CuS04r 
K2S04,6H2O,  which  are  deposited  in  large,  monosymmetric  crystals 
from  the  mixed  solutions  (see  Zinc  sulphate). 

The  Sulphides  of  Copper.  —  Cuprous  sulphide  Cu2S  occurs  in 
nature  in  rhombic  crystals  of  a  gray,  metallic  appearance.  It  is 
the  sulphide  formed  by  direct  union  of  the  elements. 

Cupric  sulphide  CuS  is  deposited  as  a  black  precipitate  when 
hydrogen  sulphide  is  led  through  a  solution  of  a  cupric  salt.  By 
cautiously  treating  copper  with  excess  of  sulphur  at  114°  it  may  be 
obtained  as  a  blue  crystalline  solid.  At  higher  temperatures  it 
gives  off  sulphur. 

Analytical  Reactions  of  Compounds  of  Copper.  —  The  ion 

of  ordinary  cupric  salts,  cupric-ion  Cu++,  is  blue,  and  that  of  cuprous 
salts,  cuprous-ion  Cu+,  is  colorless.  Cuprous  solutions,  however,  are 
easily  oxidized  by  the  air  and  become  blue.  In  solutions  containing 
cupric-ion,  hydrogen  sulphide  precipitates  cupric  sulphide,  even  in 
presence  of  acids  (p.  421).  Bases  throw  down  the  blue  hydroxide, 
and  carbonates  precipitate  a  green  basic  salt  (p.  744) .  Potassium  f  er- 
rocyanide  gives  the  brown,  gelatinous  cupric  f  errocyanide :  2Cu.SO4  + 
K4.Fe(CN)6  <^  Cu2.Fe(CN)6 1  +  2K2SO4.  A  very  characteristic  test 
is  the  formation  of  the  deep-blue  Cu(NH3)444"  ion  with  excess  of 
ammonium  hydroxide.  This  solution,  because  of  the  very  slight 
concentration  of  Cu++,  gives  a  precipitate  with  hydrogen  sulphide 
only.  Solutions  of  complex  cuprous  cyanides  such  as  potassium 
cuprocyanide  K.Cu(CN)2  are  colorless,  and  do  not  respond  to  any 
of  the  above  tests.  With  microcosmic  salt  or  borax  (pp.  560,  640), 
copper  compounds  form  a  bead  which  is  green  in  the  oxidizing  part 
of  the  flame  and  becomes  red  and  opaque  (liberation  of  copper)  in 
the  reducing  flame. 

Electrotyping.  —  When  plates  of  platinum,  connected  with  a 
battery,  are  immersed  in  cupric  sulphate  solution,  copper  is  deposited 


COPPER,   SILVER,   GOLD 


747 


on  the  cathode  (negative  pole).  The  sulphate-ion  S04~  migrates 
(p.  347)  towards  the  anode  (positive  pole)  and  there  produces  sul- 
phuric acid  and  oxygen  (p.  344) .  If,  however,  the  anode  is  made  of 
copper,  the  SO4~  migrates,  but  is  not  discharged.  Instead,  copper 
goes  into  solution  (Fig.  147)  as  Cu"1"1",  in  amount  equal  to  that  de- 
posited on  the  other  pole.  Thus,  the  only  changes  are,  (1)  an  in- 
crease in  concentration  of  cupric  sulphate 
round  the  positive  pole  (anode),  and  (2) 
a  transfer  of  copper  from  the  copper 
anode  to  the  cathode  (see  below). 

A  copper  electrotype  of  a  medal  (or 
a  page  of  type)  is  made  by  first  prepar- 
ing a  cast  of  the  medal  in  plaster  of 
Paris,  gutta  percha,  or  wax.  The  sur- 
face of  the  cast  is  then  rubbed  with 
graphite,  to  render  it  a  conductor,  and 
the  cast  is  then  used  as  the  cathode  in  a  cell  with  a  copper  anode, 
like  that  just  described.  The  deposit  of  copper,  when  heavy  enough, 
is  stripped  off.  In  making  book  plates,  the  cast  is  made  with  wax, 
+  and  the  copper  electrotype  is  strengthened  and 
thickened  by  filling  the  back  with  melted  lead.* 


Fia.  147. 


Copper  Refining.  —  The  tenacity,  duc- 
tility, and  conductivity  of  copper  are  seriously 
affected  by  small  amounts  of  impurities,  such 
as  cuprous  oxide  or  sulphide,  which  are  soluble 
in  the  molten  metal.  Arsenic  amounting  to 
0.03  per  cent  lowers  the  conductance  about 
14  per  cent.  There  are  also  silver  and  gold 
in  smelter  copper.  Hence,  a  large  propor- 
tion of  the  copper  on  the  market  is  purified 
by  electrolysis.  The  principle  is  the  same 
as  that  used  in  electrotyping.  Thin  sheets 
of  copper  form  the  cathodes,  and  thick  plates 
of  crude  copper  the  anodes.  These  are  suspended  alternately 
and  close  together  in  large  troughs,  lined  with  lead,  and  filled  with 
cupric  sulphate  solution  (Fig.  148,  diagrammatic,  view  from  above). 
The  cathodes  are  all  connected  with  the  negative  wire  of  the  dynamo, 
and  the  anodes  with  the  positive  one.  The  Cu++  is  attracted  to  the 

*  For  newspapers,  a  plate  is  made  from  the  cast  of  the  type  more  quickly 
by  means  of  melted  stereotype  metal  (lead,  antimony,  tin;  82  :  15  :  3). 


FIG.  148. 


748  INORGANIC  CHEMISTRY 

cathodes  and  is  deposited  upon  them.  The  SC>4=  migrates  towards 
the  anodes,  where  copper  from  the  thick  plate  forms  ions  Cu"^  in 
equivalent  amount.  The  stock  of  cupric  sulphate  thus  remains  the 
same,  and  the  liquid  is  stirred  to  keep  the  sulphate  from  accumulat- 
ing close  to  the  anodes.  The  practical  effect  of  the  electrolysis  is 
to  carry  copper  across  from  one  plate  to  the  other.  The  cathodes 
are  removed  from  time  to  time,  and  the  deposit  of  copper  is  stripped 
from  their  surface.  Fresh  anodes  are  substituted  when  the  old  ones 
are  eaten  away.  Since  there  is  no  final  decomposition  of  any  cupric 
sulphate,  the  only  electrical  energy  required  is  that  necessary  to  over- 
come the  friction  of  the  moving  ions.  Hence,  a  very  small  difference 
in  potential  (less  than  0.5  volts)  is  sufficient  (see  p.  799). 

The  less  active  metals,  which  are  mixed  with  the  copper  in  the 
anode,  are  not  ionized,  because  there  is  plenty  of  the  more  active 
copper  to  carry  the  current.  These  metals,  and  traces  of  sulphides, 
therefore,  fall  to  the  bottom  of  the  vat  as  a  sludge.  Zinc  and  other 
metals  more  active  than  copper,  however,  are  ionized.  Conversely, 
at  the  cathode,  the  copper,  being  the  least  active  metal  present 
in  ionic  form,  is  alone  deposited.  There  is  no  tendency  to  dis- 
charge zinc  or  hydrogen,  for  example,  so  long  as  there  are  plenty 
of  the  more  easily  discharged  copper  ions  available  (see  p.  799). 
In  this  way,  copper,  99.8  per  cent  pure,  is  obtained,  gold  and  silver 
are  recovered  from  the  sludge,  and  the  bath  liquid  is  removed 
from  time  to  time  for  purification  from  the  more  active  metals  it 
acquires.  In  1914,  in  the  United  States  alone,  264,825  ounces  of  gold 
and  nearly  fifteen  million  ounces  of  silver  were  obtained  in  this  way. 

SILVER  AG 

Chemical  Relations  of  the  Element.  —  This  element  presents 
a  curious  assortment  of  chemical  properties.  It  differs  from  copper 
in  having  a  strongly  basic  oxide,  in  giving  salts  with  active  acids 
which  are  not  hydrolyzed  by  water,  and  in  forming  neutral  rather 
than  basic  salts.  In  these  respects  it  approaches  the  metals  of  the 
alkalies  and  alkaline  earths.  It  resembles  copper  in  entering  into 
complex  compounds,  and  in  giving  insoluble  halides  like  the  cuprous 
halides.  It  differs  from  both  copper  and  the  metals  of  the  alkalies, 
and  resembles  gold  and  platinum,  in  that  its  oxide  is  easily  decom- 
posed by  heat,  with  formation  of  the  free  metal,  and  in  the  low  posi- 
tion it  occupies  in  the  electromotive  series  and  the  consequent 
slight  chemical  activity  of  the  free  metal. 


COPPER,   SILVER,   GOLD  749 

The  salts  are  always  represented  by  the  simplest  formula,  AgCl, 
etc.,  although  in  organic  solvents  greater  tendencies  to  polymeriza- 
tion are  observed  than  in  the  case  of  the  cuprous  compounds  (p.  735). 

Occurrence.  —  Native  silver,  sometimes  found  in  large  masses, 
although  more  usually  scattered  through  a  rocky  matrix,  contains 
varying  amounts  of  gold  and  copper.  Native  copper  always  contains 
dissolved  silver.  Sulphide  of  silver  Ag2S  occurs  alone  and  dissolved 
in  galenite  PbS,  with  which  it  is  isomorphous.  Smaller  amounts  of 
the  metal  are  obtained  from  pyrargyriteAg3SbS3andproustite  Ag3AsS3, 
which  are  silver  sulphantimonite  and  sulpharsenite,  respectively,  and 
from  horn-silver  AgCl.  The  chief  supplies  come  from  California, 
Australia,  and  Mexico. 

Metallurgy.  —  The  silver  contained  free,  or  as  sulphide,  in  ores 
of  copper  and  lead,  is  found  in  the  free  state  dissolved  in  the  metals 
extracted  from  these  ores,  and  is  secured  by  refining  them.  In  the 
electrolytic  refining  of  copper,  silver  is  obtained  from  the  sludge 
deposited  in  the  baths  (p.  747).  The  proportion  present  in  lead  is 
usually  small.  Formerly  the  Pattinson  desilverizing  process  was 
largely  employed.  In  it  the  metallic  lead  is  melted  in  iron  vessels, 
and  the  crystals  of  lead,  deposited  as  the  metal  slowly  loses  heat, 
are  raked  out.  These  consist  at  first  of  pure  lead  (cf.  p.  199).  When 
the  remaining  liquid  becomes  saturated  with  silver  it  begins  to 
deposit  lead  and  silver  together.  At  this  point  the  residue  is  placed 
in  a  hollow,  lined  with  bone-ash,  forming  part  of  a  reverberatory 
furnace  (Fig.  143,  p.  686),  and  heated  strongly  while  a  blast  of 
air  passes  over  its  surface.  In  this  process,  called  "  cupellation, " 
the  lead  is  converted  into  litharge  PbO,  which,  driven  by  the  air, 
flows  in  molten  condition  over  the  edge  of  the  cupel.  When  the 
last  trace  of  lead  is  gone,  the  shining  surface  of  the  pure  silver  " flashes" 
into  view  (cf.  p.  659).  Parke's  process,  which  has  superseded  the 
above,  takes  advantage  of  the  fact  that  molten  zinc  and  lead  are 
practically  insoluble  in  one  another,  while  silver  is  much  more  solu- 
ble in  zinc  than  in  lead.  Lead  dissolves  1.6  per  cent  of  zinc,  and 
zinc  1.2  per  cent  of  lead.  The  principle  is  the  same  as  in  the  re- 
moval of  iodine  from  water  by  ether  (p.  275).  The  lead  is  melted 
and  thoroughly  mixed  by  machinery  with  a  small  proportion  of 
zinc.  After  a  short  time  the  zinc  floats  to  the  top,  carrying  with  it  in 
solution  almost  all  of  the  silver,  and  solidifies  at  a  temperature  at 
which  the  lead  is  still  molten.  The  zinc-silver  alloy,  largely  a  com- 


750  INORGANIC   CHEMISTRY 

pound  Ag2Zri5,  is  skimmed  off,  and  heated  moderately  in  a  furnace 
to  permit  the  adhering  lead  to  drain  away.  The  zinc  is  finally 
distilled  off  in  clay  retorts,  and  the  lead  remaining  with  the  silver 
is  removed  by  cupellation. 

The  gold  which  goes  with  the  silver  in  Parke's  process  is  sepa- 
rated electrolytically  (p.  747).  Plates  of  the  silver-gold  alloy  form 
the  anode,  and  silver  nitrate  solution  the  vat-liquid.  The  silver, 
being  the  more  active  metal,  is  ionized  and  deposited  on  the  cathode, 
while  the  gold  collects  as  a  powder  in  a  bag  surrounding  the  anode. 

Ores  of  silver  which  do  not  contain  much  or  any  lead  are  often 
smelted  with  lead  ores,  and  the  product  is  treated  as  described  above, 
but  many  other  processes  are  in  use.  In  Mexico  the  "patio"  proc- 
ess has  been  in  use  since  1557.  The  sulphide  is  converted  into 
chloride  by  the  action  of  cupric  chloride.  Metallic  mercury  dis- 
places the  silver:  AgCl  +  Hg  — >  HgCl  +  Ag,  and,  being  present  in 
excess,  dissolves  it.  The  treatment  occupies  several  weeks,  and 
much  mercury  is  consumed.  The  amalgam  is  finally  secured  by 
"  washing, "  and  the  mercury  is  separated  from  the  silver  by  distilla- 
tion. 

During  the  first  half  of  the  nineteenth  century  the  total  world's 
output  averaged  only  643  tons  per  year.  Up  to  1870  a  gram  of 
gold  could  buy  15.5  g.  of  silver.  Now  that  the  production  has  reached 
about  8000  tons,  the  same  amount  of  gold  purchases  about  40  g. 
The  chief  sources  (1911)  are  Mexico  2460  tons,  United  States  1880, 
Canada  1018,  Europe  525. 

Physical  Properties.  —  Pure  silver  is  almost  perfectly  white. 
It  melts  at  960°.  Its  density  is  10.5.  Its  ductility  is  so  great  that 
wires  can  be  drawn  of  such  fineness  that  2  kilometers  weigh  only 
about  1  g.  In  the  molten  condition  it  absorbs  mechanically  about 
twenty-two  times  its  own  volume  of  oxygen,  but  gives  up  almost  all 
of  this  as  it  solidifies.  Fantastically  irregular  masses  result  from  the 
"sprouting"  or  "spitting"  which  accompanies  the  escape  of  the 
gas. 

By  addition  of  ferrous  citrate  to  silver  nitrate,  a  red  solution  and 
lilac  precipitate  of  free  silver  can  be  made.  The  latter,  after  washing 
with  ammonium  nitrate  solution,  gives  a  red  colloidal  suspension 
(p.  621)  in  water.  It  is  a  negatively  charged  colloid,  and  is  coagulated 
by  bivalent  positive  ions.  Colloidal  (cf.  p.  621)  silver  showing  a 
variety  of  colors,  due  to  different  degrees  of  dispersion,  has  been 
prepared  by  Gary  Lea.  Colloidal  suspensions  of  metals  are  formed 


COPPER,   SILVER,   GOLD  751 

also  by  passing  an  electrical  discharge  between  wires  of  silver,  gold, 
or  platinum  held  under  water. 

Silver  is  alloyed  with  copper  to  render  it  harder.  The  silver  coin- 
age of  the  United  States  and  the  continent  of  Europe  has  a  "fine- 
ness of  900"  (900  parts  of  silver  in  1000),  that  of  Great  Britain  925. 
Silver  ornaments  have  a  fineness  of  800  or  more.  A  superficial  layer 
of  almost  pure-white  silver  is  produced  by  heating  the  object  in  the 
air  and  dissolving  out  the  cupric  oxide  thus  formed  with  dilute 
sulphuric  acid.  The  surface  of  the  products,  if  not  subsequently  bur- 
nished, is  "frosted."  "Oxidized  silver"  is  made  by  dipping  objects 
made  of  the  metal  in  a  solution  of  potassium-hydrogen  sulphide, 
whereby  a  thin  film  of  silver  sulphide  is  produced. 

Chemical  Properties.  —  Silver,  when  cold,  is  oxidized  by  ozone, 
but  not  by  oxygen  (see  silver  oxide).  It  does  not  ordinarily  displace 
hydrogen  from  aqueous  solutions  of  acids,  but  its  tendency  to  form  the 
sulphide  is  so  great  that  it  decomposes  hydrogen  sulphide  and  alkali 
sulphides  (cf.  p.  441).  Hence,  sulphur  compounds  in  the  air  tarnish 
the  surface,  producing  Ag2S,  as  do  also  eggs,  secretions  from  the 
skin  (proteins,  p.  628),  and  vulcanized  rubber.  It  also  displaces 
hydrogen  when  boiled  with  concentrated  hydriodic  acid,  giving 
AgHI2  (cf.  Aqua  regia,  p.  537).  Silver  interacts  with  cold  nitric 
acid  and  with  hot,  concentrated  sulphuric  acid,  giving  the  nitrate 
or  sulphate  of  silver  and  oxides  of  nitrogen  or  of  sulphur  (p.  535). 
Since  its  hydroxide  has  no  tendency  to  behave  as  an  acid,  alkalies, 
whether  in  solution  or  fused,  have  no  action  upon  silver.  Hence  alka- 
line substances  are  heated  in  vessels  of  this  metal  or  of  iron,  rather 
than  in  vessels  of  platinum  (q.v.),  because  platinum  is  attacked  by 
alkaline  materials. 

The  Halides  of  Silver.  —  The  chloride  AgCl,  bromide  AgBr, 
and  iodide  Agl  are  formed  as  curdy  precipitates  when  a  salt  of 
silver  is  added  to  a  solution  containing  the  appropriate  halide  ion. 
The  first  is  white,  and  melts  at  about  457°.  The  second  and  third 
are  very  pale-yellow  and  yellow  respectively.  The  insolubility  in 
water,  which  is  very  great,  increases  in  the  above  order.  The  iodide, 
after  melting,  solidifies  and  forms  quadratic  crystals,  which,  as  they 
cool,  pass  at  146°  (transition  point)  into  a  different  physical  variety 
(hexagonal)  with  evolution  of  heat  (cf.  pp.  679,  738). 

When  exposed  to  light,  the  chloride  becomes  first  violet  (colloidal 
silver,  dispersed  in  the  AgCl)  and  finally  brown,  chlorine  being 


752  INORGANIC  CHEMISTRY 

liberated.  The  bromide  and  iodide  behave  similarly.  Solid  silver 
chloride  absorbs  ammonia,  forming  with  a  low  pressure  2AgCl,- 
3NH3,  and  with  a  higher  pressure  of  the  gas  AgCl,3NH3,  the  former 
with  a  tension  of  93  mm.,  and  the  latter  with  a  tension  of  about  one 
atmosphere  of  ammonia  at  20°  (cf.  p.  154).  The  bromide  forms  no 
compound  in  this  way,  but  the  iodide  yields  2AgI,NH3. 

In  consequence  of  the  progressive  insolubility,  a  cold  solution  of  a 
bromide  will  convert  the  precipitate  of  silver  chloride  into  bromide, 
and  a  soluble  iodide  will  similarly  transform  the  bromide  or  the 
chloride  into  iodide  (cf.  p.  405). 

Silver  fluoride  AgF  may  be  made  by  treating  the  oxide  or  car- 
bonate with  hydrofluoric  acid:  H2F2  +  Ag2O  -»  2AgF  +  H20.  The 
salt  is  very  soluble  and  deliquescent. 

Complex  Compounds  of  Silver.  —  Silver  chloride  interacts 
easily  with  excess  of  ammonium  hydroxide,  giving  the  complex 
cation  Ag(NH3)2+.  Under  certain  conditions  octahedral  crystals 
(Fig.  63,  p.  172)  of  AgCl  are  deposited  from  the  .solution,  and,  under 
other  conditions,  crystals  of  the  composition  2AgCl,3NH3.  The 
bromide,  which  interacts  less  readily,  gives  the  same  complex  ion. 
The  iodide  hardly  interacts  at  all.  Ammonio-argentic-ion  Ag(NH3)2+, 
in  solutions  of  concentrations  such  as  are  dommoiily  used  (0.1  N  to  N), 
gives  about  the  same  concentration  of  silver-ion  Ag+  as  does  the 
bromide,  and  much  more  than  the  highly  insoluble  iodide  (cf.  p.  742) . 
Hence  the  latter  interacts  very  slightly  with  ammonium  hydrox- 
ide, and  can  be  precipitated  in  ammoniacal  solution.  All  three 
of  the  insoluble  halides  dissolve  in  solutions  of  potassium  cyanide  and 
of  sodium  thiosulphate,  as  do  also  all  the  other  insoluble  silver  salts. 
Usually  an  equivalent  amount  of  the  cyanide  or  thiosulphate  suffices, 
but  for  solution  of  the  sulphide  an  excess  is  required.  With  the 
cyanide,  double  decomposition  gives  first  the  insoluble  silver  cyanide 
AgCN  which  then  interacts  forming  the  soluble  potassium  argenticya- 
nide  K.Ag(CN)2.  The  thiosulphate  gives  a  solution  from  which  crystals 
of  a  complex  salt  Na3.Ag(S203)2  are  obtained.  The  complex  anion 
in  the  solution  appears  to  be  Ag(S203)2-.  Since  the  iodide  dissolves 
in  the  thiosulphate  with  considerable  difficulty,  we  should  infer 
that  the  complex  thiosulphate  anion  gives  about  the  same  concentra- 
tion of  argentic-ion  as  does  the  iodide.  An  independent  method  of 
measuring  the  concentrations  of  argentic-ion  in  all  the  solutions  places 
the  compounds  in  the  order  of  diminishing  _ability  to  give  argentic- 
Jon  thus,  AgCl?  Ag(NB3)2+,  AgBr,  Ag(S203)2=  Agl,  Ag(CN)r,  Ag2S, 


COPPER,   SILVER,   GOLD  753 

and  confirms  the  above  inferences  (see  Concentration  cells).  The 
more  active  metals,  like  zinc  and  copper,  displace  silver  from  all 
solutions,  whether  the  solutions  contain  simple  or  complex  salts. 

Oxide  of  Silver.  —  When  sodium  hydroxide  is  added  to  a  solu- 
tion of  a  salt  of  silver,  a  pale-brown  precipitate  is  obtained,  which, 
after  being  freed  from  water,  is  found  to  be  Ag2O.  We  should 
expect  to  obtain  the  hydroxide  AgOH  in  this  fashion,  but  it  appears 
to  be  unstable.  The  aqueous  solution  of  argentic  oxide,  however, 
is  distinctly  alkaline,  and  presumably  therefore  does  contain  the 
hydroxide:  2 AgOH  +±  Ag2O  +  H2O.  Silver  oxide  is  formed  by  boil- 
ing silver  chloride  with  caustic  potash.  Since  the  oxide  is  much  more 
soluble  than  the  chloride  (see  table),  we  should  expect  the  reverse  of  the 
above  action  to  be  the  normal  one.  Here,  however,  the  excess  of 
potassium  hydroxide  (hydroxide-ion)  represses  the  ionization  of  the 
silver  hydroxide  and  reverses  the  relations  in  regard  to  solubility 
(cf.  p.  698). 

Argentic  oxide  is  an  active  basic  oxide,  and  all  the  salts  of  silver 
are  derived  from  it.  When  moist,  it  absorbs  carbon  dioxide  from 
the  air.  Its  solutions  show  concentrations  of  hydroxide-ion  smaller, 
it  is  true,  than  equimolar  solutions  of  the  active  bases,  but  con- 
siderably greater  than  similar  solutions  of  ammonium  hydroxide 
(p.  367).  With  ammonium  hydroxide  it  forms  the  soluble  ammonio- 
argentic  hydroxide  Ag(NH3)2.OH,  which  is  as  active  a  base  as  is 
potassium  hydroxide.  The  solution,  when  allowed  to  evaporate, 
deposits  black  crystals  of  an  explosive  substance  whose  composition 
has  not  been  determined.  This  is  " fulminating  silver"  (not  to  be 
confused  with  fulminate  of  silver  Ag.ONC).  When  the  oxide  is 
heated,  it  gives  off  oxygen,  leaving  metallic  silver.  The  action  is 
reversible  (G.  N.  Lewis)  and  the  dissociation  pressure  of  the  oxygen 
is  20.5  atmospheres  at  302°.  At  a  higher  pressure  than  this  (at  302°), 
therefore,  oxygen  will  combine  with  silver. 

Silver  Nitrate  AgNOz.  —  This  salt  is  obtained  by  treating  silver 
with  aqueous  nitric  acid: 

3Ag  +  4HN03  ->  3AgN03  +  NO  +  2H20. 

From  the  solution,  colorless  rhombic  crystals  (Fig.  14,  p.  19)  isomor- 
phous  with  those  of  potassium  nitrate  (Fig.  71,  p.  173)  are  deposited. 
These  melt  at  208.6°.  Thin  sticks  made  by  casting  (lunar*  caustic) 

*  (Lat.)  luna  (the  moon),  the  alchemical  name  for  silver. 


754  INORGANIC  CHEMISTRY 

are  used  to  cauterize  sores,  because  the  substance  combines  with 
proteins  to  form  insoluble  compounds.  When  commercial  silver,  con- 
taining copper,  is  used  to  make  silver  nitrate,  the  solution  is  evapo- 
rated to  dryness  and  heated  at  250°  until  the  nitrate  of  copper  has 
all  been  decomposed.  At  this  temperature  the  silver  salt  is  unaffected 
and,  when  cool,  can  be  separated  from  the  insoluble  cupric  oxide  by 
extraction  with  water. 

The  aqueous  solution  is  neutral.  The  pure  salt  is  not  affected  by 
light,  but  when  deposited  on  cloth,  on  the  skin  of  the  fingers,  or  on 
the  mouth  of  the  reagent  bottle,  it  is  converted  into  the  chloride,  and 
from  this,  in  turn,  silver  is  liberated.  For  this  reason  it  is  an  ingredi- 
ent in  marking-inks.  The  dry  compound  combines  with  ammonia, 
giving  AgN03,3NH3.  When  added  to  an  aqueous  solution,  am- 
monium hydroxide  produces,  first  a  faint  precipitation  of  the  oxide, 
and  then  the  soluble  complex  salt  Ag(NH3)2.N03. 

Other  Salts  of  Silver.  —  Silver  peroxidate  Ag202  (cf.  p.  318)  is 
formed  by  the  action  of  ozone  on  silver.  In  the  electrolysis  of  silver 
nitrate  a  deposit  of  shining  black  crystals  (3Ag2O2,AgN03,2O)  is 
formed  on  the  positive  electrode.  When  they  are  boiled  with  water, 
oxygen  is  given  off,  silver  nitrate  dissolves,  and  the  peroxidate  is 
deposited  as  a  gray  powder.  Silver  carbonate,  the  neutral  salt 
Ag2C03,  and  not  a  basic  carbonate,  is  precipitated  from  solutions  of 
salts  of  silver  by  soluble  carbonates.  It  is  slightly  yellow  in  color. 
With  water  it  gives  a  faint  alkaline  reaction  and,  like  calcium  carbon- 
ate, is  soluble  in  excess  of  carbonic  acid  (p.  705).  When  heated,  the 
carbonate  decomposes,  leaving  metallic  silver.  Other  compounds  of 
silver,  for  example,  the  chloride,  when  heated  in  a  crucible  with 
sodium  carbonate  give  this  salt  by  double  decomposition,  and  hence 
are  finally  reduced  to  a  button  of  metallic  silver.  The  sulphate 
Ag2SO4  is  made  by  the  action  of  concentrated  sulphuric  acid  on  the 
metal.  It  is  not  very  soluble  in  water,  and  crystallizes  in  rhombic 
prisms  isomorphous  with  anhydrous  sodium  sulphate.  When  it  is 
mixed  with  a  solution  of  aluminium  sulphate  (q.v.),  octahedral  crys- 
tals of  silver-alum  Ag2SO4,Al2(S04)3,24H2O  are  obtained.  Silver 
sulphide  Ag2S  is  precipitated  by  hydrogen  sulphide  from  solutions 
of  all  silver  compounds,  whether  free  acids  are  present  or  not,  and 
irrespective  of  the  form  in  which  the  silver  is  combined.  Excess  of 
potassium  cyanide,  however,  prevents  its  precipitation  from  the 
argenticyanide.  The  sulphide  is  formed  by  the  action  of  metallic 
silver  on  alkaline  hydrosulphides,  and  this  interaction  forms  the 


COPPER,   SILVER,   GOLD  755 

basis  of  the  "hepar"  test  for  sulphur  (p.  441).  Silver  orthophos- 
phate  Ag3PO4  (yellow),  arsenate  Ag3As04  (brown),  and  chromate 
Ag2CrO4  (crimson)  are  produced  by  precipitation,  and  their  dis- 
tinctive colors  enable  us  to  use  silver  nitrate  in  analysis  as  a  reagent 
for  identifying  the  acid  radicals. 

Electroplating.  —  The  process  is  similar  to  the  electro-deposi- 
tion of  copper  (p.  747).  The  article  to  be  plated  is  cleaned  with 
extreme  care  and  attached  to  the  negative  wire.  A  plate  of  silver 
forms  the  positive  electrode  and,  since  simple  salts  of  silver  do  not 
give  coherent  deposits,  the  bath  is  a  solution  of  potassium  argenti- 
cyanide.  The  potassium-ion  K+  migrates  to  the  negative  wire,  and 
since  potassium  requires  a  much  greater  E.M.F.  for  its  liberation  than 
does  silver,  silver  is  there  deposited  from  the  trace  of  silver-ion 
given  by  the  complex  silver  ions  in  the  neighborhood: 

Ag(CN)2-  ±+  Ag+  +  2CN-        Ag+  +  0  -+  Ag. 

The  potassium  cyanide  remains  in  solution.  At  the  positive  elec- 
trode silver  goes  into  solution  in  equivalent  amount  giving  argentic- 
ion,  and  the  above  equations  are  reversed. 

Mirrors  are  silvered  through  the  reduction  of  ammonio-silver 
nitrate  by  organic  compounds  such  as  potassium-sodium  tartrate 
(Rochelle  salt),  glycerine,  formaldehyde  CH2O  (formalin),  or  grape 
sugar:  4AgOH  -f  CH2O  -+  3H2O  -f  4Ag  j  +  C02.  On  a  small 
scale,  dilute  silver  nitrate  is  mixed  with  ammonium  hydroxide  until 
the  solution  is  clear,  and  then  a  little  caustic  potash,  a  few  more 
drops  of  ammonia,  and  finally  a  very  little  glycerine  are  added.  A 
watch-glass  floated  on  this  mixture  quickly  acquires  a  deposit  of 
silver.  The  film  of  silver  is  washed,  dried,  and  varnished. 

Photography.  —  Bromo-gelatine  dry  plates  are  made  by  pre- 
paring an  emulsion  of  gelatine  to  which  silver  nitrate  and  a  slight 
excess  of  ammonium  bromide  have  been  added.  After  the  emulsion 
has  been  kept  warm  until  the  precipitate  of  silver  bromide  has  coagu- 
lated into  small  granules  (" ripening"),  it  is  allowed  to  solidify.  It  is 
then  cut  up,  and  the  ammonium  nitrate  is  washed  out  with  water. 
After  drying  and  remelting,  the  emulsion  is  finally  applied  to  plates 
of  glass  or  films  of  celluloid.  The  excess  of  ammonium  bromide  and 
the  ripening  both  increase  the  subsequent  sensitiveness  of  the  plates. 
Scott  Archer  (1850)  first  suspended  the  silver  halide  in  a  jelly  (col- 
lodion). 


756  INORGANIC  CHEMISTRY 

After  exposure,  often  for  only  a  fraction  of  a  second,  there  is  no 
visible  alteration  in  the  film.  The  image  is  developed.  Chemically, 
this  consists  in  reducing  the  silver  bromide  to  metallic  silver  by  means 
of  reducing  agents.  While  the  whole  of  the  halide  upon  the  plate  is 
reducible,  if  the  reducing  agent  is  kept  upon  it  for  a  sufficient  length 
of  time,  the  parts  reached  by  the  light  are  affected  first,  and  with  a 
speed  proportional  to  the  intensity  of  the  illumination  undergone  by 
each  part.  The  reducing  agent  is  poured  off  when  sufficient  "con- 
trast" between  the  parts  variously  illuminated  has  been  attained. 
Development  was  discovered  by  Talbot  (1837),  or  rather  by  his  cat, 
which  upset  an  extract  of  nut-galls  on  some  half-exposed  papers  cov- 
ered with  silver  chloride.  The  unreduced  silver  bromide  is  then  dis- 
solved out  with  sodium  thiosulphate  (" hyposulphite  of  soda"  or 
"hypo"),  first  used  by  Sir  John  Hershell  (1840),  and  the  silver  image 
is  thus  saved  from  obliteration  by  the  silver  that  would  be  deposited 
if  the  plate  were  to  be  brought  into  the  light  without  this  treatment 
(fixing) .  The  result  is  a  "  negative, "  as  the  parts  brightest  in  the  obj  ec t 
are  now  opaque,  and  the  darkest  parts  of  the  object  are  transparent. 

The  gelatine  is  the  sensitizing  substance,  and  promotes  the  disso- 
ciation of  the  silver  bromide:  2AgBr  ^±  2Ag  +  Br2,  which  is  a 
reversible  action,  by  combining  with  the  bromine.  Potassium  bro- 
mide is  added  to  the  developer  if  the  plate  has  been  over-exposed. 
This  restrains  the  development,  by  rendering  the  silver  bromide  less 
soluble  (cf.  p.  698). 

A  common  developer  is  the  potassium  salt  of  hydroquinone 
C6H4(OH)2,  which  gives  quinone  C6H402: 

2AgBr  +  (KO)2C6H4  ->  2Ag  +  2KBr  +  C6H402. 

An  alkaline  solution  of  the  sodium  salt  of  pyrogallic  acid  is  often 
used. 

In  printing,  the  light  and  dark  are  again  reversed,  the  denser  parts 
of  the  negative  protecting  the  compounds  on  the  paper  below  it  from 
action,  and  leaving  them  white.  Either  "bromide"  papers  (such  as 
velox,  invented  by  Baekeland),  which  require  only  brief  exposure 
and  are  developed  like  the  plate,  are  used,  or  silver  chloride  is  the 
sensitive  substance,  and  prolonged  exposure  to  light  is  allowed  to 
liberate  the  proper  amount  of  silver.  The  operation  of  fixing  is 
performed  as  before.  In  toning  chloride  papers,  a  solution  of  sodium 
chloraurate  is  employed.  A  portion  of  the  silver  dissolves,  displacing 
gold  (p.  404),  which  is  deposited  in  its  place: 

NaAuCU  +  3Ag  ->  NaCl  +  3AgCl  +  Au. 


COPPER,   SILVER,   GOLD  757 

The  thin  film  of  gold  gives  a  richer  color  to  the  print.     In  platinum 
toning,  potassium  chloroplatinite  K2PtCl4  is  similarly  used. 

Many  other  actions  are  utilized  in  photography.  Thus,  ferric 
oxalate  is  reduced  by  light  to  ferrous  oxalate:  Fe2(C204)3  —  >  2FeC204 
+  2CO2.  When  paper  coated  with  a  solution  of  the  former,  or  with  a 
mixture  of  ferric  chloride  and  ammonium  oxalate,  is  used  for  printing, 
the  pale-yellow  ferric  salt  loses  its  color  where  it  has  been  turned  into 
the  ferrous  salt.  If  the  paper  is  then  dipped  in  a  solution  of  ferri- 
cyanide  of  potassium  K3.Fe(CN)e  the  ferrous  salt  precipitates  the 
insoluble  and  deep-blue  ferrous  ferricyanide  Fe3[Fe(CN)6]2,  while  the 
unchanged  ferric  salt  simply  gives  a  soluble  brown  substance,  which 
can  be  washed  out.  For  regular  blue  prints  ammonium-ferric  citrate 
is  employed  instead  of  the  oxalate.  If  the  above  paper,  after  print- 
ing, is  dipped  in  potassium  chloroplatinite  (or  has  been  coated  with 
this  salt  at  the  same  time  that  it  received  the  ferric  oxalate),  and  is 
then  dipped  in  potassium  oxalate  solution,  the  latter  dissolves  the  in- 
soluble ferrous  oxalate,  and  the  potassium-ferrous  oxalate  reduces  the 
platinum  compound,  giving  a  platinum  print: 


6FeC2O4  +  3K2PtCl4  ->  2Fe2(C2O4)3  +  2FeCl3  +  3Pt  +  6KC1. 

We  have  already  seen  (p.  581)  that  light  of  short  wave-length  — 
blue  and  violet  —  has  the  greatest  effect  upon  silver  halides.  The 
time,  in  seconds,  required  for  equal  effects  is  approximately:  violet 
15,  blue  29,  green  37,  yellow  330,  red  600.  Hence  objects  showing 
to  the  eye  a  variety  of  colors  are  entirely  misrepresented,  as  regards 
the  relative  brightness  of  their  parts,  by  photography.  Now  the  im- 
portant fact,  in  this  connection  is,  that  only  that  part  of  the  light 
which  is  absorbed  in  traversing  the  film,  and  not  that  which  is  scat- 
tered or  transmitted,  can  be  used  for  chemical  change.  Hence,  dip- 
ping plates  in  solutions  of  substances  capable  of  absorbing  yellow  and 
red  radiations  causes  them  to  absorb  more  of  the  energy  of  these 
photographically  weakest  radiations,  and  to  give  greater  chemical 
action  in  response  to  them.  This  partially  restores  the  balance. 
Such  plates  are  called  orthochromatic,  and  are  backed  with  sub- 
stances like  eosin  (used  also  in  making  red  ink)  or  cyanine. 

Analytical  Reactions  of  Silver  Compounds.  —  Argentic-ion 
Ag+  is  colorless.  Many  of  its  compounds  are  insoluble,  the  pre- 
cipitation of  the  chloride,  which  is  insoluble  in  dilute  acids,  being 
used  as  a  test.  Mercurous  chloride  and  lead  chloride  are  also  white 
and  insoluble,  but,  in  ammonium  hydroxide,  silver  chloride  dissolves, 


758  INORGANIC   CHEMISTRY 

mercurous  chloride  (q.v.)  turns  black,  and  lead  chloride,  which  is 
also  soluble  in  hot  water,  is  not  altered  in  color.  With  excess  of 
ammonium  hydroxide,  silver  salts  give  the  complex  cation  Ag(NH3)2+ 
and,  from  solutions  containing  silver  in  this  form,  only  the  iodide  and 
sulphide  can  be  precipitated.  Sodium  thiosulphate  and  potassium 
cyanide  dissolve  all  silver  salts,  giving  salts  of  complex  acids  with 
silver  in  the  anion  (p.  752).  Zinc  displaces  silver  from  all  forms  of 
combination. 

GOLD  Au 

Chemical  Relations  of  the  Element.  —  This  element  forms 
two  very  incomplete  series  of  compounds  corresponding,  respectively, 
to  aurous  and  auric  oxides,  Au2O  and  Au203.  The  former  is  a  feebly 
basic  oxide,  the  latter  acid-forming.  No  simple  salts  with  oxygen 
acids  are  stable.  All  the  compounds  of  gold  are  easily  decomposed 
by  heat  with  liberation  of  the  metal.  All  other  common  metals 
displace  gold  from  solutions  of  its  compounds  (p.  404).  Mild  re- 
ducing agents  likewise  liberate  gold.  The  element  enters  into  many 
complex  anions  (p.  649). 

Occurrence  and  Metallurgy.  —  Gold  is  found  chiefly  in  the 
free  condition  disseminated  in  veins  of  quartz,  or  mixed  with  alluvial 
sand.  Small  quantities  are  found  also  in  sulphide  ores  of  iron,  lead, 
and  copper.  Telluride  of  gold  (sylvanite),  in  which  silver  takes  the 
place  of  part  of  the  gold  [Au^\g]Te2*,  is  found  in  Colorado.  This 
mineral  when  heated  loses  its  tellurium,  and  gold,  alloyed  with  silver, 
remains. 

From  the  alluvial  deposits,  gold  is  usually  separated  by  washing 
in  a  cradle  (densities,  gold  19.32,  rock  about  2.6),  as  in  the  Klondyke. 
Quartz  veins,  which  in  the  Transvaal  Colony  reach  a  thickness  of  a 
meter  and  carry  an  average  of  18  g.  of  gold  per  ton,  are  mined,  and 
the  material  is  pulverized  with  stamping  machinery.  About  55  per 
cent  of  the  gold  is  then  separated  by  allowing  the  powdered  rock  to 
be  carried  by  a  stream  of  water  over  copper  plates  amalgamated  with 
mercury.  The  gold  dissolves  in  the  latter,  and  is  secured  by  removal 
and  distillation  of  the  amalgam.  The  finer  particles  contained  in 
the  sludge  which  runs  off  (" tailings")  are  extracted  by  adding  a 

*  Amongst  minerals,  mixed  crystals  of  isomorphous  salts  are  so  com- 
monly found  that  formulae  like  the  above  are  constantly  used  by  mineralogists. 
[Au,Ag]Te2  indicates  a  mixture  in  varying  proportions  of  the  isomorphous  tellu- 
rides  AuTe2  and  AgTej. 


COPPER,   SILVER,   GOLD  759 

dilute  solution  of  sodium  cyanide  (MacArthur-Forest  process)  and 
exposing  the  mixture  to  the  air.  Oxidation  and  simultaneous  inter- 
action with  the  cyanide  give  sodium  aurocyanide  Na.Au(CN)2. 
Hydrogen  peroxide,  which  is  formed  in  many  oxidations  by  free 
oxygen,  is  produced  also: 

2Au  +  4NaCN  +  2H2O  +  02  ->  2NaAu(CN)2  +  2NaOH  +  H2O2, 
2Au  +  4NaCN  +  H202  -»  2NaAu(CN)2  +  2NaOH. 

From  this  solution  the  gold  is  isolated,  either  by  electrolysis,  in 
which  a  plate  of  lead  forms  the  cathode  (and  is  subsequently  cupelled : 
Siemens-Halske  process),  or  in  the  form  of  a  purple  powder  by  pre- 
cipitation with  zinc.  The  same  cyanide  is  used  for  another  batch. 

The  gold  is  separated  from  ores  containing  silver,  copper,  lead,  and 
other  metals,  and  various  methods  of  refining,  mainly  electrolytic, 
are  used. 

The  world's  production  of  gold  during  the  first  half  of  the  nine- 
teenth century  averaged  27  tons  annually.  In  1897  it  was  363  tons, 
and  in  1899,  472.6  tons.  This  rapid  increase  in  the  supply  of  gold 
(which  is  our  standard  of  value)  has  made  it  relatively  cheaper,  and 
other  articles  more  expensive.  In  1913  the  total  production  was 
680  tons,  of  which  the  Transvaal  gave  40  per  cent,  the  United  States 
20  per  cent,  and  Australia  12  per  cent,  Russia  6  per  cent,  Mexico 
4  per  cent,  Canada  4  per  cent,  India  nearly  3  per  cent. 

Properties  of  the  Metal.  —  Gold  is  yellow  in  color,  and  is  the 
most  malleable  and  ductile  of  all  the  metals.  It  melts  at  1063°.  Its 
density  is  19.32.  To  give  it  greater  hardness  it  is  alloyed  with  cop- 
per, the  proportion  of  gold  being  defined  in  "  carats."  Pure  gold 
is  "24-carat."  British  sovereigns  are  22-carat  and  contain  &  of 
copper.  American,  French,  and  German  coins  are  21.6-carat,  or 
90  per  cent  gold.  Silver  takes  the  place  of  copper  in  Australian  sov- 
ereigns. 

Gold  is  not  affected  by  free  oxygen  or  by  hydrogen  sulphide.  It 
does  not  displace  hydrogen  from  dilute  acids,  nor  does  it  interact  with 
nitric  or  sulphuric  acids  or  any  oxygen  acids,  except  selenic  acid.  It 
combines,  however,  with  free  chlorine  (and  bromine),  and  it  therefore 
interacts  with  a  mixture  of  nitric  and  hydrochloric  acids  (aqua  regia), 
which  gives  off  this  gas  (p.  537),  giving  chlorauric  acid  H.AuCL* 
(=  HCljAuCla).  This  happens,  not  because  aqua  regia  is  more 
active  than  are  any  of  the  substances  it  contains,  but  because  it 
furnishes  both  the  chlorine  and  the  chloride-ion  Cl~~  required  to  pro- 


760  INORGANIC  CHEMISTRY 

duce  the  exceedingly  stable  (little  dissociated)  anion  AuCLf.  Chlo- 
rine-water (C12,H+,C1-,C1O-)  dissolves  it  also,  for  the  same  reason. 
Gold  is  the  least  active  of  the  familiar  metals. 

Compounds  with  the  Halogens.  —  Chlorauric  acid,  formed  as 
above,  is  deposited  in  yellow,  deliquescent  crystals  of  H.AuCLt,  4H20. 
The  yellow  sodium  chloraurate  NaAuCl4,2H2O,  obtained  by  neu- 
tralization of  the  acid,  is  used  in  photography  (p.  756).  The  acid 
gives  up  hydrogen  chloride  when  heated  very  gently,  leaving  the  red, 
crystalline  auric  chloride  AuCl3.  The  tendency  to  form  complex 
compounds  is  such,  however,  that  when  dissolved  in  water  free  from 
hydrochloric  acid,  this  salt  gives  H2.AuCl3O.  Red  crystals  of 
H2AuCl3O,2H2O  are  deposited  by  the  solution.  When  auric  chloride 
is  heated  to  180°  aurous  chloride  AuCl  and  chlorine  are  formed.  This 
salt  is  a  white  powder.  It  is  insoluble  in  water,  but  in  boiling  water 
is  converted  quickly  into  auric  chloride  and  free  gold :  3 AuCl  — » 
AuCl3  +  2Au.  When  potassium  iodide  is  added  to  a  solution  of 
chlorauric  acid,  or  to  sodium  chloraurate,  the  yellow  aurous  iodide 
is  precipitated: 

NaAuCU  +  3KI  -> NaCl  +  Aul  J,  +  I2  +  3KC1. 

The  action  is  like  that  on  cupric  salts  (p.  741),  and  for  a  similar 
reason,  namely,  that  auric  iodide  is  not  stable. 

Other  Compounds.  —  When  caustic  alkalies  are  added  to 
chlorauric  acid,  or  to  sodium  chloraurate,  auric  hydroxide  Au(OH)3 
is  precipitated.  This  substance  is  an  acid,  and  interacts  with  excess 
of  the  base,  forming  aurates.  These  are  derived  from  met-auric  acid 
(Au(OH)s  —  H2O  =  HAu02),  as,  for  example,  potassium  aurate 
K.AuC>2,3H2O.  This  salt  interacts  by  double  decomposition,  giving, 
for  instance,  with  silver  nitrate,  the  insoluble  silver  aurate  AgAu02. 
Its  solution  is  alkaline  in  reaction,  a  fact  which  shows  that  auric  acid 
is  a  weak  acid  (cf.  p.  648).  Auric  oxide  Au203  is  a  brown,  and 
aurous  oxide  Au2O  is  a  violet  powder.  With  hydrochloric  acid  the 
latter  gives  chlorauric  acid  and  free  gold. 

On  account  of  its  reducing  action,  hydrogen  sulphide  precipitates 
from  chlorauric  acid  a  dark-brown  mixture  containing  much  aurous 
sulphide  Au2S  and  free  sulphur,  as  well  as  some  auric  sulphide  Au2S3. 
The  sulphides  interact  with  alkali  sulphides,  giving  complex  sulphau- 
rites  and  sulphaurates,  such  as  K3AuS2  (=  3K2S,Au2S)  and  KAuS2 


COPPER,   SILVER,  GOLD  761 

(=  K2S,Au2S3),  which  are  soluble  (p.  649,  and  see  Tin,  Arsenic,  and 
Antimony) . 

The  aurocyanides,  like  K.Au(CN)2  (=  KCN,AuCN),  and  the 
auricyanides,  like  K.Au(CN)4  (=  KCN,Au(CN)3),  are  formed  by  the 
action  of  potassium  cyanide  on  aurous  and  auric  compounds,  re- 
spectively. They  are  colorless  and  soluble.  Their  solutions  are 
used  as  baths,  in  conjunction  with  a  gold  anode,  for  electrogilding. 

Analytical  Reactions  of  Gold.  —  The  metallic  "streak,"  pro- 
duced by  rubbing  the  metal  on  touchstone  (Lydian  stone,  a  black 
basalt),  is  not  easily  removed  by  nitric  acid  of  sp.  gr.  1.36  (57.5  per 
cent).  The  fineness  of  the  gold  can  be  determined  by  comparing  the 
effect  of  the  acid  with  that  on  streaks  from  pieces  of  gold  of  known 
fineness. 

In  assaying,  the  material  containing  the  gold  is  heated  with  borax 
and  lead  in  a  small  crucible  (cupel)  of  bone-ash.  The  lead  and  copper 
are  oxidized,  and  the  oxides  are  absorbed  by  the  cupel,  leaving  a  drop 
of  molten  alloy  of  gold  and  silver.  The  cold  button  is  flattened  by 
hammering  and  rolling,  and  treated  with  nitric  acid  to  remove  the 
silver.  The  gold,  which  remains  unattacked,  is  washed,  fused  again, 
and  weighed.  The  acid  will  not  interact  with  the  silver  and  remove  it 
completely  if  the  quantity  of  gold  exceeds  25  per  cent.  When  the 
proportion  of  gold  is  greater  than  this,  a  suitable  amount  of  pure 
silver  is  fused  with  the  alloy  ("quartation"). 

Exercises.  —  1.  How  much  copper  will  be  deposited  per  hour 
on  each  sq.  cm.  of  an  electrode  immersed  in  cupric  sulphate  solution 
when  the  current  density  is  \  ampere  per  sq.  cm.  (p.  357)?  How 
much  copper  would  be  obtained  under  the  same  conditions  from  a 
cuprous  salt? 

2.  Write  equations  for  the  interactions,  (a)  of  salt  water  and 
oxygen  with  copper  (p.  738),  (6)  of  ferrous  oxide  and  sand  (p.  737), 
(c)  of  verdigris,  arsenious  acid,  and  acetic  acid  (p.  745). 

3.  Write  the  formulae  of  the  basic  chloride,  nitrate,  carbonate, 
and  sulphate  of  copper  as  if  these  substances  were  composed  of  the 
normal  salt,  the  oxide  and  water  (p.  738). 

4.  What  may  be  the  formula  of  the  compound  of  cupric  hydroxide 
and  sodium  tartrate  (p.  744)? 

5.  Can  you  develop  any  relation  between  the  facts  that  solutions 
of  cupric  salts  are  acid  in  reaction  and  that  they  give  basic  carbon- 
ates by  precipitation? 


762  INORGANIC   CHEMISTRY 

6.  Formulate  the  action  of  potassium  cyanide  in  dissolving  cupric 
hydroxide  and  cuprous  sulphide,  assuming  that  potassium  cupro- 
cyanide is  formed. 

7.  How  should  you  set  about  making  cupric  orthophosphate  (in 
solution),  ammonium  cuprocyanide,  and  lead  cuprocyanide? 

8.  Write  the  formulae  of  some  of  the  double  salts  analogous  to 
potassium-cupric  sulphate  (p.  746). 

9.  What  chemical  reagents  are  present  in  a  Bunsen  flame?     If 
borax  beads  were  made  in  the  oxidizing  flame  with  cupric  chloride, 
cuprous  bromide,  and  cupric  sulphate,  severally,  what  actions  would 
take  place? 

10.  If  the  solubility  ratio  of  silver  in  zinc  and  in  lead  were  1000  : 1, 
and  2  per  cent  of  zinc  were  used,  what  proportion  of  the  total  silver 
would  be  secured  by  Parke's  method? 

11.  Which  is  more  stable,  (a)  silver  sulphate  or  cupric  sulphate, 
(6)  silver  nitrate  or  cupric  nitrate?     To  what  salts  are  the  silver 
compounds  in  this  respect  more  closely  allied? 

12.  Write  the  equations  for  the  interaction  of,  (a)  silver  and  con- 
centrated sulphuric  acid,  (6)  silver  chloride  and  sodium  carbonate, 
when  heated  strongly,  (c)  sodium  thiosulphate  and  silver  bromide, 
(d)  potassium  ferricyanide  and  ferrous  oxalate. 

13.  What  reagents  should  you  use  to  precipitate  the  phosphate, 
arsenate,  and  chromate  of  silver? 

14.  Write  the  equations  for  the  interactions  of,  (a)  gold  and  selenic 
acid,  in  which  selenious  acid  is  formed,  (6)  potassium  hydroxide  and 
auric  hydroxide,  (c)  potassium  cyanide  and  sodium  chloraurate. 

15.  In  what  respects  are  the  elements  of  this  family  distinctly 
metallic,  and  in  what  respects  are  they  allied  to  the  non-metals 
(p.  645)? 

16.  Collect  all  the  evidence  tending  to  show  that  the  cuprous  com- 
pounds are  more  stable  than  the  cupric. 

17.  Describe  in  terms  of  the  categories  used  by  the  phase  rule  the 
systems,  (a)  cupric  nitrate  and  water  at  24.5°  and  (6)  silver  iodide  at 
146°. 

18.  Make  a  classified  list  of  the  methods  by  which  cupric  com- 
pounds are  transformed  into  cuprous,  and  vice  versa. 

19.  Of  which  metals  should  it  be  possible  to  obtain  colloidal 
suspensions  in  water,  and  of  which  not  (p.  404)  ? 


CHAPTER  XXXVII 

GLUCINUM,   MAGNESIUM,   ZINC,   CADMIUM,   MERCURY.     THE 
RECOGNITION   OF   CATIONS   IN   QUALITATIVE  ANALYSIS 

The  Chemical  Relations  of  the  Family.  —  The  remaining 
elements  of  the  third  column  of  the  periodic  table,  namely,  glucinum 
or  beryllium  (Gl,  or  Be,  at.  wt.  9.1),  magnesium  (Mg,  at.  wt.  24.32), 
zinc  (Zn,  at.  wt.  65.37),  cadmium  (Cd,  at.  wt.  112.4),  and  mercury 
(Hg,  at.  wt.  200.6),  although  all  bivalent,  do  not  form  a  coherent 
family.  Glucinum  and  magnesium  resemble  zinc  and  cadmium,  and 
differ  from  the  calcium  family,  in  that  the  sulphates  are  soluble,  the 
hydroxides  easily  lose  water  leaving  the  oxides,  the  chlorides  are  com- 
paratively volatile,  and  the  metals  are  not  rapidly  rusted  in  the  air 
and  do  not  easily  displace  hydrogen  from  water.  They  resemble  the 
calcium  family,  and  differ  from  zinc  and  cadmium,  in  that  the  sul- 
phides are  hydrolyzed  by  water,  the  oxides  are  not  reduced  by  heating 
with  carbon,  complex  cations  are  not  formed  with  ammonia,  and  the 
metals  do  not  enter  into  complex  anioiis.  But  glucinum  differs  from 
magnesium  and  resembles  zinc  in  that  its  hydroxide  is  acidic  as  well 
as  basic.  This  is  not  unnatural,  since  in  the  periodic  system  it  lies 
between  lithium,  a  metal,  and  boron,  a  non-metal.  Mercury  is  the 
only  member  of  the  group  that  forms  two  series  of  compounds.  These 
are  derived  (p.  488)  from  the  oxides  HgO  and  Hg2O.  Mercury  ap- 
proaches the  noble  metals  in  the  ease  with  which  its  oxide  is  decom- 
posed by  heating,  and  in  the  position  of  the  free  element  in  the 
electromotive  series. 

The  vapor  densities  of  zinc,  cadmium,  and  mercury  show  the 
vapors  of  these  three  metals  to  be  monatomic. 

The  compounds  of  the  metals  of  this  family  give  no  color  to  the 
borax  bead. 

GLUCINTJM  GL 

Chemical  Relations  of  the  Element.  —  Glucinum  (or  beryl- 
lium) is  bivalent  in  all  its  compounds.  Its  oxide  and  hydroxide  are 
basic,  and  are  also  feebly  acidic  towards  active  bases  (see  Zinc  hydrox- 
ide). On  account  of  this  fact  and  the  extreme  ease  with  which  its 

763 


764  INORGANIC  CHEMISTRY 

carbonate  gives  up  carbon  dioxide,  in  both  of  which  respects  it  re- 
sembles aluminium,  it  was  first  thought  to  be  trivalent.  This  made 
its  atomic  weight  13.6,  the  amount  combining  with  one  chemical  unit 
of  chlorine  being  4.55.  In  the  periodic  system,  however,  there  was  a 
space  for  a  bivalent  element  with  the  atomic  weight  9.1  (  =  2  X  4.55) 
between  lithium  and  boron,  and  none  for  a  trivalent  element.  Later 
(1884)  Nilson  and  Pettersson  determined  the  vapor  density  of  the 
chloride  and  of  certain  organic  compounds  of  the  element,  and  found 
only  9.1  parts  of  glucinum  in  the  molar  weights  of  the  compounds. 
The  element  derives  its  name  from  the  sweet  taste  of  its  salts  (Gk. 
vs,  sweet). 


The  Metal  and  its  Compounds.  —  Glucinum  occurs  in  beryl, 
a  metasilicate  of  glucinum  and  aluminium  Al2Gl3(SiO3)6.  Specimens 
of  beryl  tinted  green  by  the  presence  of  a  little  silicate  of  chromium 
are  known  as  emeralds.  The  metal  may  be  obtained  by  electrolysis 
of  the  easily  fusible  double  fluoride  G1F2,2KF.  In  powdered  form  it 
burns  when  heated  in  the  air.  It  displaces  hydrogen  from  dilute 
acids,  and  also,  when  heated,  from  caustic  potash  :  Gl  +  2KOH  —  > 
K2GK)2  +  H2.  The  oxide  interacts  with  acids  and  with  strong  bases. 
The  salts  give  no  color  to  the  Bunsen  flame. 

MAGNESIUM  MG 

Chemical  Relations  of  the  Element.  —  Magnesium  is  biva- 
lent in  all  its  compounds.  The  oxide  and  hydroxide  are  basic  ex- 
clusively. The  element  does  not  enter  into  complex  cations  or  anions. 

Occurrence.  —  Magnesium  carbonate  occurs  alone  as  magne- 
site,  and  in  a  double  salt  with  calcium  carbonate  MgCO3,CaCO3  as 
dolomite.  The  sulphate  and  chloride  are  found  as  hydrates  and  as 
constituents  of  double  salts  (see  below)  in  the  Stassfurt  deposits. 
Silicates  are  also  common.  Olivine  is  the  orthosilicate  Mg2SiO4. 
Serpentine  is  a  hydrated  disilicate,  [Mg,Fe]3Si2O7,2H2O,  as  is  also 
meerschaum.  Asbestos  is  an  anhydrous  silicate.  Talc  (soapstone) 
is  H2Mg3(Si03)4.  The  element  derives  its  name  from  Magnesia,  a 
town  in  Asia  Minor. 

The  Metal.  —  Magnesium  is  manufactured  by  electrolysis  of 
dehydrated  and  fused  carnallite  MgCl2,KCl,6H2O.  The  iron  cruci- 
ble in  which  the  material  is  melted  forms  the  cathode,  and  a  rod  of 


GLUCINUM,   MAGNESIUM,  ZINC,  CADMIUM,   MERCURY      765 

carbon  the  anode.  The  metal  is  silver-white,  and  when  heated  can 
be  pressed  into  wire  and  rolled  into  ribbon  (m.-p.  651°,  b.-p.  1100°). 
Commercial  specimens  of  the  latter  often  contain  zinc. 

Chemically  the  metal  is  less  active  than  are  the  metals  of  the  alka- 
line earths.  It  slowly  becomes  coated  with  a  layer  of  a  basic  car- 
bonate, which  is  non-coherent  and  so  does  not  protect  the  metal 
from  rusting  completely.  It  displaces  hydrogen  slowly  from  boiling 
water,  and  does  so  rapidly  if  some  ammonium  chloride  is  added 
to  interact  with  the  layer  of  the  hydroxide,  and,  of  course,  rapidly 
from  cold,  dilute  acids.  Magnesium  burns  in  air  with  a  white  light, 
rich  in  rays  of  short  wave-length  such  as  act  upon  photographic 
plates  (p.  757).  The  ash  contains  the  nitride  Mg3N2,  as  well  as  the 
oxide.  The  former  interacts  with  water  to  give  ammonia  (p.  514). 
When  the  metal  is  heated  with  the  oxides  of  boron,  of  silicon,  and  of 
many  of  the  metals,  it  combines  with  the  oxygen  and  liberates  the 
other  element. 

Powdered  magnesium  is  used  in  pyrotechny  and,  with  potassium 
chlorate  (10  :  17),  in  making  flash-light  powder  for  use  in  photog- 
raphy. 

Magnesium  Chloride  MgO2,6Jf2O.  —  This  salt  occurs  in  salt 
deposits.  It  is  a  highly  deliquescent  compound,  obtained  also  by 
evaporating  an  aqueous  solution,  and  as  carnallite  MgCl2,KCl,6H20. 
The  latter  is  an  important  source  of  potassium  chloride  (p.  664),  and 
almost  all  the  magnesium  chloride  combined  with  it  is  thrown  away. 
When  the  hexahydrate  is  heated,  a  part  of  the  chloride  is  hydrolyzed, 
some  magnesium  oxide  remaining,  and  some  hydrogen  chloride  being 
given  off.  Sea-water  cannot  be  used  in  ships'  boilers  because  of  the 
hydrochloric  acid  liberated  by  the  magnesium  chloride  which  the 
water  contains.  The  salt  forms  a  double  chloride  with  ammonium 
chloride  MgCl2,NH4Cl,6H20  which  is  isomorphous  with  carnallite, 
and  this  salt  can  be  dehydrated  without  hydrolysis  of  the  chloride. 
Afterwards  the  ammonium  chloride  can  be  volatilized  (p.  520).  To 
utilize  natural  magnesium  chloride,  the  manufacture  of  chlorine  from 
it,  by  passing  air  and  steam  over  the  salt  at  a  high  temperature,  has 
been  attempted: 

4MgCl2  +  2H20  +  02  ->  4MgO  +  4HC1  +  2C12. 

The  Oxide  and  Hydroxide.  —  Magnesium  oxide  MgO  is  made 
by  heating  the  carbonate,  and  is  known  as  calcined  magnesia.  It  is 
a  white,  highly  infusible  powder,  and  is  used  for  lining  electric  fur- 


766  INORGANIC   CHEMISTRY 

naces  and  making  crucibles.  It  combines  slowly  with  water  to  form 
the  hydroxide  Mg(OH)2. 

The  hydroxide  is  found  in  nature  as  brucite.  It  is  also  precipi- 
tated from  solutions  of  magnesium  salts  by  alkalies.  It  is  very 
slightly  soluble  in  water,  much  less  so  than  calcium  hydroxide,  but 
more  so  than  are  the  hydroxides  of  zinc  and  the  other  heavy  metals. 
The  solution  has  a  faint  alkaline  reaction.  When  magnesium  chlo- 
ride is  added  to  the  moist  hydroxide,  a  hydrated  basic  chloride, 
(Mg(OH)2)2,(MgCl2)1/,(H20)z  is  formed.  The  mixture,  to  which 
sawdust  is  sometimes  added,  is  used  as  a  plaster-finish  in  building. 

Magnesium  hydroxide  is  not  precipitated  by  ammonium  hydroxide 
when  ammonium  salts  are  present  also.  The  ammonium  salts,  being 
highly  ionized  and  giving  a  high  concentration  of  ammonium-ion 
NH4+,  repress  the  ionization  of  the  feebly  ionized  ammonium  hydrox- 
ide, and  so  reduce  the  concentration  of  hydroxide-ion  which  it  fur- 
nishes. With  the  ordinary  concentration  of  Mg4"1",  therefore,  the 
amount  of  hydroxide-ion  existing  in  presence  of  excess  of  a  salt  of 
ammonium  is  too  small  to  bring  the  ion  product  [Mg++]  X  [OH~]2 
up  to  the  value  required  for  precipitation  (p.  698).  Conversely,  mag- 
nesium hydroxide  interacts  with  solutions  of  ammonium  salts  and 
passes  into  solution: 

Mg(OH)2(solid)  *±  Mg(OH)2  T±  Mg++  +  20H~ 
2NH4C1    ; 


In  presence  of  excess  of  ammonium  chloride,  the  OH~  combines  with 
NH4+  to  form  molecular  ammonium  hydroxide,  and  the  equilibria  in 
the  upper  line  are  displaced  forwards  to  generate  a  further  supply  of 
the  OH~.  With  sufficiently  great  concentration  of  the  ammonium 
chloride,  all  the  magnesium  hydroxide  may  thus  dissolve;  with  only  a 
small  excess,  a  condition  of  equilibrium  with  solid  magnesium  hydrox- 
ide is  reached.  The  whole  case  is  analogous  to  the  interaction  of 
acids  with  insoluble  salts  (p.  713).  Magnesium  oxide  also  dissolves 
in  salts  of  ammonium.  It  gives  first  the  hydroxide  by  interaction 
with  the  water. 

Magnesium  Carbonates.  —  The  normal  carbonate  MgCOs  (mag- 
nesite)  is  found  in  nature.  Only  hydrated  basic  carbonates  are 
formed  by  precipitation,  and  their  composition  varies  with  the 
conditions.  The  carbonate  manufactured  in  large  amounts  and 
sold  as  magnesia  alba  is  approximately  Mg4(OH)2(C03)3.3H20.  It 


GLUCINUM,   MAGNESIUM,   ZINC,   CADMIUM,   MERCURY     767 

is  used  in  medicine  and  as  a  cosmetic.  The  carbonates  are  not  pre- 
cipitated in  the  presence  of  ammonium  salts,  and  interact  with  such 
salts  in  the  same  way  as  does  the  hydroxide. 

Magnesium  Sulphate  MgSOi.  —  The  common  heptahydrate 
MgSO4,7H2O  crystallizes  from  cold  water  in  rhombic  prisms,  and  is 
called  Epsom  salts.  At  0°  a  dodecahydrate  appears.  The  hepta- 
hydrate is  efflorescent,  and  loses  its  water  by  stages  and  with  decreas- 
ing aqueous  tension.  The  monohydrate,  found  in  the  salt  layers  as 
kieserite  MgS04,H2O,  has  a  very  low  aqueous  tension,  and  is  not 
rapidly  dehydrated  except  above  200°.  The  hepta-  and  monohy- 
drates  present  a  striking  case  of  difference  in  solubility  in  two  forms 
of  one  salt,  the  former  giving  at  15°  a  solution  containing  33.8  g.  of 
the  sulphate  in  100  g.  of  water,  while  the  latter  is  almost  insoluble. 
Magnesium  sulphate  is  used  in  the  manufacture  of  sodium  and 
potassium  sulphates,  and  is  employed  also  for  " loading"  cotton 
goods,  and  as  a  purgative. 

Magnesium  Sulphide  MgS.  —  The  sulphide  may  be  formed 
by  heating  the  metal  with  sulphur.  It  is  insoluble  in  water,  but  is 
decomposed  and  gives,  finally,  hydrogen  sulphide  and  magnesium 
hydroxide : 

2MgS  +  2H20  t^Mg(SH)2  +  Mg(OH)2, 
Mg(SH)2  +  2H20  <±  Mg(OH)2  J  +  2H2S. 

The  hydrolysis  is  more  complete  than  in  the  case  of  calcium  sulphide, 
and  eliminates  all  the  hydrogen  sulphide,  because  magnesium  hydrox- 
ide is  much  more  insoluble  than  is  calcium  hydroxide,  and  so  there  is 
little  reverse  interaction  tending  to  reproduce  the  soluble  hydro- 
sulphide  Mg(SH)2. 

Phosphates  of  Magnesium.  —  The  only  phosphate  of  impor- 
tance is  ammonium-magnesium  orthophosphate  NH4MgPO4,6H20, 
which  appears  as  a  crystalline  precipitate  when  sodium  phosphate  and 
ammonium  hydroxide  (and  chloride,  p.  766)  are  mixed  with  a  solution 
of  a  magnesium  salt.  This  compound  is  insoluble  in  water  containing 
ammonium  hydroxide,  and  is  used  in  quantitative  analysis  for  esti- 
mating both  magnesium  and  phosphoric  acid.  Before  being  weighed 
the  precipitate  is  ignited,  and  is  thus  converted  into  the  anhydrous 
pyrophosphate  of  magnesium  Mg2P2O7.  The  salt  NH4MgAsO4,6H2O 
has  similar  properties,  and  is  used  for  estimating  arsenic  acid. 


768  INORGANIC  CHEMISTRY 

Analytical  Reactions  of  Magnesium  Compounds. — The  mag- 
nesium ion  Mg++  is  colorless  and  bivalent.  It  does  not  enter  into 
complex  ions.  Soluble  carbonates  precipitate  basic  carbonates  of 
magnesium,  but  not  when  ammonium  salts  are  present.  The  latter 
limitation  distinguishes  compounds  of  magnesium  from  those  of  the 
calcium  family.  Potassium  hydroxide  precipitates  the  hydroxide  of 
magnesium,  except  when  salts  of  ammonium  are  present.  The 
mixed  phosphate  of  ammonium  and  magnesium,  in  presence  of 
ammonium  hydroxide,  is  the  least  soluble  salt. 

ZINC  ZN 

Chemical  Relations  of  the  Element.  —  Zinc  is  bivalent  in  all 
its  compounds.  Of  these  there  are  two  sets,  —  the  more  numerous 
and  important  one  in  which  zinc  is  the  positive  radical  (Zn.SO4, 
Zn.Cl2,  etc.),  and  a  less  numerous  set,  the  zincates,  in  which  zinc  is 
in  the  negative  radical  (Na.HZnO2,  etc.).  Both  sets  of  salts  are 
hydrolyzed  by  water,  as  the  hydroxide  is  feeble  whether  it  is  con- 
sidered as  an  acid  or  as  a  base.  The  element  also  enters  into  com- 
plex cations  and  anions.  The  salts  are  all  poisonous. 

Occurrence  and  Extraction  from  the  Ores.  —  The  chief 
sources  of  zinc  are  calamine  Zn2SiO4,H20,  smithsonite  ZnCO3,  zinc- 
blende  (Ger.  blenden,  to  dazzle)  or  sphalerite  ZnS,  franklinite 
Zn.(FeO2)2,  and  zincite  ZnO.  The  red  color  of  the  last  is  due  to 
the  presence  of  manganese. 

The  ores  are  first  concentrated  by  froth  flotation  (p.  737).  They 
are  then  converted  into  oxide  —  the  carbonate  by  ignition,  and  the 
sulphide  by  roasting.  The  sulphur  dioxide  is  used  to  make  sulphuric 
acid.  A  mixture  of  the  oxide  with  coal  is  then  distilled  in  earthen- 
ware retorts  at  1300-1400°,  the  zinc  condensing  in  earthenware 
receivers,  while  carbon  monoxide  burns  at  a  small  opening: 

2ZnS  +  3O2  ->  2ZnO  +  2S02, 
ZnO  +  C  ->  CO  +  Zn. 

At  first  zinc  dust,  a  mixture  of  zinc  and  zinc  oxide,  collects  in  the 
receiver,  and  afterwards  liquid  zinc.  The  product,  which  is  cast  in 
blocks,  is  called  spelter.  It  contains  small  amounts  of  lead,  arsenic, 
iron,  gallium,  and  cadmium,  because  the  sulphides  of  these  metals  are 
almost  invariably  present  in  zinc-blende.  The  amounts  of  spelter 
manufactured  (1913)  were  United  States  346,700  short  tons,  Germany 
312,000,  Belgium  218,000,  Great  Britain  65,000. 


GLUCINUM,   MAGNESIUM,  ZINC,   CADMIUM,  MERCURY     769 

Properties  and  Uses  of  the  Metal.  —  Zinc  is  a  bluish-white, 
crystalline  metal.  When  cold  it  is  brittle,  but  at  120-150°  it  can  be 
rolled  into  sheets  between  heated  rollers  and  then  retains  its  pliability 
when  cold.  At  200-300°  the  metal  becomes  once  more  brittle,  at 
419.4°  it  melts,  and  at  925°  it  boils.  The  vapor  density  at  1740°  is 
2.64  (air  =  1),  and  the  molecular  weight,  therefore,  2.38  X  28.955 
(p.  233)  or  68.9.  The  gas  is  thus  monatomic. 

The  metal  burns  in  air  with  a  greenish  flame,  giving  zinc  oxide. 
When  cold  it  is  not  affected  by  dry  air,  but  in  moist  air  it  is  oxidized, 
and  becomes  covered  with  a  firmly  adhering,  non-porous  layer  of 
basic  carbonate  which  protects  it  from  further  action.  The  metal 
displaces  hydrogen  from  dilute  acids,  but  with  pure  specimens  the 
action  almost  ceases  in  consequence  of  the  formation  of  a  layer  of 
condensed  hydrogen  on  the  surface.  Contact  with  a  less  electro- 
positive metal,  such  as  lead,  iron,  copper,  or  platinum,  enables  the 
action  to  go  on,  because  the  hydrogen  is  then  liberated  at  the  surface 
of  the  other  metal  (see  Electromotive  chemistry).  Crude  zinc  con- 
tains lead  and  iron  and  is  therefore  more  active  than  pure  zinc. 
Zinc  also  attacks  boiling  alkalies,  giving  the  soluble  zincate  (see 
below) :  KOH  +  Zn  +  H2O  -->  KHZn02  +  H2.  The  action  on  am- 
monium hydroxide  is  slower  and  different  in  n.ature: 

Zn  +  2H20  +  4NH3  -»  Zn(NH3)4.(OH)2  +  H2, 

a  complex  cation  being  formed. 

Sheet  zinc,  in  consequence  of  its  lightness  (density  7),  is  used  in 
preference  to  lead  (density  11.5)  for  roofs,  gutters,  and  architectural 
ornaments.  Galvanized  iron  is  made  by  dipping  sheet  iron,  cleaned 
with  sulphuric  acid  or  the  sand  blast,  into  molten  zinc.  The  latter, 
being  more  active  (p.  404),  is  rusted  instead  of  the  iron,  but  the  rust- 
ing is  very  slight.  Objects  of  iron,  cleaned  and  baked  in  zinc  dust, 
also  acquire  a  coating  of  zinc  (sherardizing) .  Zinc  is  used  also  in 
batteries  and  for  making  alloys  (p.  644) .  It  mixes  in  all  proportions 
with  tin,  copper,  and  antimony,  but  with  lead  (p.  749)  and  with 
bismuth  separation  into  two  layers  occurs,  each  metal  dissolving  only 
a  little  of  the  other.  The  two  different  modes  of  behavior  resemble 
these  of  alcohol  and  water  (p.  179)  and  ether  and  water  (p.  180), 
respectively. 

Zinc  Chloride  ZnClz.  —  This  salt  is  usually  manufactured  by 
treating  zinc  with  excess  of  hydrochloric  acid,  evaporating  the  solu- 
tion to  dryness,  and  fusing  the  residue.  When  hydrochloric  acid  is 


770  INORGANIC  CHEMISTRY 

thus  present,  the  chloride  ZnCl2  is  obtained.  Evaporation  of  the 
pure  aqueous  solution,  which  is.  acid  in  reaction,  results  in  consider- 
able hydrolysis  and  formation  of  much  of  the  basic  chloride  Zn2OCl2: 

ZnCl2  +  H20  *±  HC1  +  Zn(OH)Cl,  (1) 

2Zn(OH)Cl  ->  Zn2OCl2  +  H20.  (2) 

The  salt  is  used  in  solid  form  as  a  caustic  and,  by  injection  of  a  solu- 
tion into  wood  (e.g.,  railway  sleepers),  as  a  poison  to  prevent  the 
growth  of  organisms  which  promote  decay.  In  both  cases  the  salt 
combines  with  proteins,  forming  solid  products.  The  aqueous  solu- 
tion, being  acid,  is  employed  also  for  dissolving  the  oxides  from 
surf  aces  which  are  to  be  soldered.  The  acid  is  reproduced  by  hydrol- 
ysis as  fast  as  it  is  used,  and  finally  the  oxychloride  remains  (equation 
1  above).  The  hot  solution  also  gelatinizes  and  dissolves  cellulose 
(cotton  or  paper),  or  probably  takes  it  into  colloidal  suspension. 
When  the  solution  is  pressed  through  an  orifice  into  alcohol,  the  cel- 
lulose is  precipitated  in  the  form  of  a  thread.  By  carbonizing  such 
threads,  carbon  filaments  for  incandescent  lamps  are  made. 

Zinc  Oxide  and  Hydroxide  and  the  Zincates.  —  The  oxide 

ZnO  is  obtained  as  a  white  powder  by  burning  zinc  or  by  heating  the 
precipitated  basic  carbonate.  It  turns  yellow  when  heated,  recover- 
ing its  whiteness  when  cold,  in  the  same  way  that  mercuric  oxide  is 
brown  whilst  hot  and  bright  red  when  cold.  It  is  used  in  making 
a  paint  —  zinc-white  or  Chinese  white  —  which  is  not  darkened  by 
hydrogen  sulphide.  It  is  also  used  as  a  filler  in  rubber  automobile 
tires.  For  filling  teeth,  dentists  sometimes  use  a  paste  made  by 
mixing  the  oxide  with  a  strong  solution  of  zinc  chloride.  It  quickly 
sets  to  a  hard  mass  of  oxychloride. 

The  hydroxide  Zn(OH)2  of  zinc  appears  as  a  white,  flocculent 
solid  when  alkalies  are  added  to  solutions  of  zinc  salts.  It  interacts 
as  a  basic  hydroxide  with  acids,  giving  salts  of  zinc: 

Zn(OH)2  +  H2S04  *=*  Zn.SO4  +  2H20. 

It  also  interacts  with  excess  of  the  alkali  employed  to  precipitate  it, 
giving  a  soluble  zincate,  such  as  potassium  zincate  KHZn02: 

H2ZnO2T  +  KOH  <=±  K.HZn02  +  H2O. 

Both  actions  are  reversible,  and  the  second  requires  a  considerable 
excess  of  alkali  for  its  completion:  in  fact,  some  of  the  zinc  hydroxide 


GLUCINUM,   MAGNESIUM,  ZINC,   CADMIUM,   MERCURY      771 

seems  to  be  simply  in  colloidal  suspension.  It  is  evident  that  zinc 
hydroxide  when  in  solution  is  ionized  both  as  an  acid  and  as  a  base: 

H+  +  HZnO2~  «=±  Zn(OH)2  (dslvd)  +±  Zn++  +  20H~ 

K 

Zn(OH)2  (solid) 

The  ionization  as  an  acid  is  less  than  that  as  a  base,  but  both  are 
small.  Addition  of  an  acid  like  sulphuric  acid,  however,  furnishes 
hydrogen-ion,  the  hydroxide  ions  combine  with  this  to  form  water, 
and  all  the  equilibria  are  displaced  to  the  right.  With  a  base,  on  the 
other  hand,  the  hydrogen-ion  is  removed  and  the  basic  ionization 
simultaneously  repressed,  so  that  the  equilibria  are  displaced  to  the 
left.  Such  a  substance  is  called  amphoteric. 

Zinc  hydroxide  interacts  with  ammonium  hydroxide,  giving  a 
soluble  complex  compound  ammonio-zinc  hydroxide  Zn(NH3)4.(OH)2. 
The  case  is  like  those  of  copper  (p.  743)  and  silver  hydroxides  (p.  753), 
and  not  like  that  of  magnesium  hydroxide  (p.  766). 

Compounds  of  zinc,  when  heated  in  the  Bunsen  flame  with  a  salt 
of  cobalt,  give  a  zincate  of  cobalt  (Rinmann's  green)  CoZn02. 

Hydrogen  sulphide  precipitates  zinc  sulphide  from  solutions  of 
zincates  and  from  solutions  containing  ammonia,  so  that  some  zinc 
ions  Zn++  are  present  in  both. 

Carbonates  of  Zinc.  —  The  normal  zinc  carbonate  ZnC03  may 
be  precipitated  by  means  of  sodium  bicarbonate: 

ZnS04  +  2NaHC03  -»  Na^SO*  +  ZnC03  +  H20  +  C02. 


The  normal  carbonate  of  sodium,  however,  gives  basic  carbonates, 
such  as  Zn2(OH)2CO3,  which,  as  in  the  case  of  magnesium  (p.  766), 
vary  in  composition  according  to  the  conditions. 

Zinc  Sulphate  ZnSO*.  —  This  salt  is  formed  by  oxidation 
when  zinc-blende  is  roasted.  It  gives  rhombic  crystals  of  the  hydrate 
ZnSO4,7H2O.  This,  and  the  corresponding  compounds  of  magne- 
sium MgSO4,7H20,  of  iron  FeS04,7H20,  and  of  several  other  bivalent 
metals,  are  all  isomorphous,  and  are  known  as  vitriols.  The  zinc 
salt  is  white  vitriol.  Like  Epsom  salts,  it  is  dehydrated  by  stages, 
the  last  molecule  of  water  being  difficult  to  remove.  It  is  used  in 
cotton-printing  and  as  an  eye-wash  (J  per  cent  solution). 

The  salt  gives  double  salts  with  potassium  or  ammonium  sul- 
phate, of  the  form  ZnS04,K2SO4,6HaO,  which  crystallize  in  the  mono- 


772  INORGANIC  CHEMISTRY 

Asymmetric  system,  and  are  isomorphous  with  each  other,  and  with 
double  salts  containing  copper  (p.  746),  mercury  (Hg11),  iron  (Fe11), 
magnesium,  and  other  bivalent  elements  in  place  of  the  zinc.  These 
compounds,  unlike  the  complex  cyanides,  are  almost  completely 
decomposed  in  dilute  solution  (cf.  p.  402). 

Zinc  Sulphide  ZnS.  —  This  compound  is  the  only  familiar  sul- 
phide which  is  white.  The  yellow  color  of  zinc-blende  is  caused  by 
the  presence  of  sulphide  of  iron.  Zinc  sulphide  is  more  soluble  in 
water  than  is  sulphide  of  copper,  and  hence  it  interacts  with  excess 
of  strong  acids,  and  passes  into  solution.  It  is  not  soluble  enough, 
however,  to  be  much  affected  by  weak  acids  like  acetic  acid.  This 
sort  of  behavior  is  shown  also  by  calcium  oxalate  (p.  713),  and  was 
discussed  fully  in  that  connection.  Zinc  sulphide  is  thus  capable  of 
being  almost  completely  precipitated  when  acetic  acid  is  present,  or 
when  hydrogen  sulphide  is  led  into  a  solution  of  the  acetate  of  zinc : 

Zn(C2H302)2  +  H2S  *±  ZnS  |  +  2HC2H3O2. 

But  when  an  active  acid  is  present,  or  is  formed,  the  sulphide  is  pre- 
cipitated very  incompletely  or  not  at  all,  the  action  being  highly 
reversible: 

ZnS04  +  H2S  *=±  ZnS  +  H2SO4. 

There  are  thus  two  ways  of  obtaining  the  sulphide  by  precipita- 
tion. A  soluble  sulphide  causes  it  to  be  thrown  down  completely 
because  no  acid  is  liberated  in  the  action : 

ZnCl2  +  (NH4)2S  <=»  ZnSJ,  +  2NH4C1. 

The  other  method  is  to  add  sodium  acetate  to  the  solution  of  the  salt, 
and  then  lead  in  hydrogen  sulphide.  The  acid,  liberated  by  the 
action  upon  the  salt,  interacts  with  the  sodium  acetate,  giving  a 
neutral  salt  of  sodium  and  acetic  acid,  and  the  zinc  sulphide  is  not 
much  affected  by  the  latter.  In  terms  of  ions,  the  hydrogen-ion, 
liberated  as  the  hydrogen  sulphide  interacts  with  the  zinc  salt,  com- 
bines with  acetate-ion  introduced  by  the  sodium  acetate,  and  givee 
the  little-ionized  acetic  acid.  For  uses,  see  lithopone  (p.  730). 

Analytical  Reactions  of  Zinc  Salts.  —  Zinc  sulphide  is  pre- 
cipitated by  the  addition  of  ammonium  sulphide  to  solutions  of  zinc 
salts  and  of  zincates.  Sodium  hydroxide  gives  the  insoluble  hydrox- 
ide, which,  however,  interacts  with  excess  of  the  alkali,  giving  the 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,  MERCURY     773 

soluble  zincate  of  sodium.     Compounds  of  zinc,  when  heated  on 
charcoal  with  cobalt  nitrate,  give  Rinmann's  green  (p.  771). 

CADMIUM  CD 

Chemical  Relations  of  the  Element.  —  This  element  is  biva- 
lent in  all  its  compounds.  Its  oxide  and  hydroxide  are  basic  exclu- 
sively, and  the  salts  are  not  hydrolyzed  by  water.  It  enters  into 
complex  compounds  having  the  ions  Cd(NH3)4++,  Cd(CN)4~,  and 


The  Metal.  —  Aside  from  the  rare  mineral  greenockite  CdS, 
cadmium  is  found  only  in  small  amounts  (about  0.5  per  cent),  as 
carbonate  and  sulphide,  in  the  corresponding  ores  of  zinc.  During 
the  reduction,  being  more  volatile  than  zinc,  it  distils  over  first 
(b.-p.  778°).  The  metal  is  white,  and  is  much  more  malleable  than 
is  zinc.  It  melts  at  320°. 

It  displaces  hydrogen  from  dilute  acids,  but  is  itself  displaced 
from  solutions  of  its  compounds  by  zinc,  since  it  is  less  electro-positive. 
It  is  used  in  making  fusible  alloys. 

Compounds  of  Cadmium.  —  The  chloride  CdCl2,2H2O  is  efflo- 
rescent and  is  not  hydrolyzed  during  dehydration  or  in  solution. 
Zinc  chloride  (p.  770)  is  deliquescent  and  is  easily  hydrolyzed.  The 
halides  are  less  ionized  than  are  the  corresponding  compounds  of 
most  other  metals.  The  iodide  CdI2,  in  particular,  seems  to  exist  in 
solution  as  Cd.CdI4,  and  the  complex  anion  gives  little  ionic  cad- 
mium. On  account  of  this  fact  even  the  sulphide  cannot  be  pre- 
cipitated completely  from  a  solution  of  the  iodide.  Conversely, 
hydriodic  acid  dissolves  the  sulphide  to  a  much  greater  extent  than 
do  other  acids  (see  below). 

The  white  hydroxide  is  made  by  precipitation,  and  interacts  as  a 
basic  hydroxide  with  acids,  but  not  at  all  with  bases.  It  dissolves 
in  ammonium  hydroxide,  however,  forming  Cd(NH3)4.(OH)2.  The 
oxide  CdO  is  a  brown  powder,  obtained  by  heating  the  hydroxide, 
carbonate,  or  nitrate,  or  by  burning  the  metal. 

The  sulphate  crystallizes  as  3CdSO4,8H2O,  and  is  not  isomor- 
phous  with  the  sulphates  of  zinc  and  magnesium.  Soluble  carbonates 
throw  down  the  normal  carbonate  CdCO3,  and  not  a  basic  carbonate. 

Hydrogen  sulphide  precipitates  the  yellow  sulphide  CdS  even 
from  acid  solutions  of  the  salts.  The  substance  is  used  as  a  pigment. 


774  INORGANIC  CHEMISTRY 

The  sulphide  of  cadmium,  however,  is  less  insoluble  in  water  than  are 
the  sulphides  of  copper  and  mercury,  and  therefore  cannot  be  precip- 
itated from  a  strongly  acid  solution  (e.g.,  HC1  >  0.3N). 

The  Solubilities  of  the  Sulphides  of  the  Metals.  —  The 

reader  will  remember  the  order  of  solubility  of  the  metallic  sul- 
phides more  easily  if  he  notes  that  it  is  practically  the  same  as  the 
order  of  activity  of  the  free  metals  (p.  404).  Thus,  the  sulphides 
down  to  that  of  aluminium  are  dissolved  by  water  (K2S  and  Na-jS), 
or  are  decomposed  by  water  (BaS,  SrS,  CaS,  MgS,  A12S3).  The 
hydroxides  formed,  being  soluble  (except  A1(OH3)),  the  whole  dis- 
solves except  in  the  case  of  A12S3.  Ferrous  sulphide  is  insoluble  in 
water,  but  is  soluble  enough  to  interact  with  (and  dissolve  in)  dilute 
acids  to  some  extent,  even  with  a  feeble  one  like  acetic  acid.  Zinc 
sulphide  requires  a  dilute  active  acid;  cadmium  sulphide  requires  a 
higher  concentration  of  an  active  acid,  as  do  also  CoS  and  NiS; 
cupric  sulphide  requires  an  oxidizing  acid  like  hot  nitric  acid;  and 
mercuric  sulphide  resists  even  this. 

The  molar  solubilities  of  the  sulphides  in  water  have  not 
been  determined  with  great  exactness,  but  are  approximately  as 
follows : 


MnS,  0.0v26 
FeS,   0.096 
ZnS,  0.0u7 

PbS,    0.0132 
CdS,  0.0i47 
Bi2S3,  O.OwS 

Ag2S,  O.Qw2 
CuS,   O.Ozol 
HgS,  0.0263 

Analytical  Reactions  of  Cadmium  Compounds.  —  The  cad- 
mium ion  Cd++  is  bivalent  and  colorless.  The  yellow  cadmium  sul- 
phide is  precipitated  by  hydrogen  sulphide,  even  from  acid  solutions 
of  the  salts.  It  is  also  precipitated  from  solutions  containing  the 
complex  cation  Cd(NH3)4++  and  the  complex  anion  Cd(CN)4=,  as,  for 
example,  from  a  solution  made  by  adding  excess  of  potassium  cyanide 
to  cadmium  chloride  solution  (K2.Cd(CN)4).  The  latter  property 
enables  cadmium  to  be  separated  from  copper  (p.  746).  The  white 
hydroxide  is  thrown  down  by  sodium  hydroxide,  and  is  not  soluble  in 
excess  of  this  reagent.  It  is  not  formed  from  solutions  containing  the 
Cd(NH3)4"H"  and  Cd(CN)4=  ions,  and  interacts  with  ammonium 
hydroxide  to  give  the  soluble  Cd(NH3)4(OH)2.  These  and  other 
precipitations  are  not  complete  when  cadmium  iodide  Cd.CdI4  is 
used. 


GLUCINUM,   MAGNESIUM,  ZINC,  CADMIUM,   MERCURY     775 

MERCURY  HG 

Chemical  Relations  of  the  Element.  —  Like  copper,  this  ele- 
ment enters  into  two  series  of  compounds,  the  mercurous  Hg1  and 
the  mercuric  Hg11,  in  which  it  is  univalent  and  bivalent,  respectively. 
The  mercurous  halides,  like  the  cuprous  halides  (and  the  argentic 
halides),  are  insoluble  in  water  and  are  decomposed  by  light.  There 
are,  however,  stable  mercurous  as  well  as  mercuric  salts  of  oxygen 
acids.  Both  of  the  oxides,  Hg2O  and  HgO,  are  basic  exclusively, 
but  in  a  feeble  degree.  The  hydroxides,  like  silver  hydroxide,  are 
not  stable,  and  lose  water,  giving  the  oxides.  The  salts  of  both  sets 
are  markedly  hydrolyzed  by.  water,  and  basic  salts  are  therefore 
common.  No  carbonate  is  known  (cf.  p.  656).  Mercury  enters 
into  the  anions  of  a  number  of  complex  salts,  such  as  HgCl4=,  HgI4=, 
Hg(CN)4=,  etc.  It  does  not  give  complex  cations  with  ammonia 
resembling  those  of  cadmium,  copper,  and  silver  (Cd(NH3)4++,  etc.), 
from  which  ammonia  is  removed  by  heating,  but  instead  forms  a 
class  of  ammono-basic  mercury  compounds  like  Hg"NH2Cl,  all  of 
which  are  insoluble.  The  mercuric  halides  and  cyanide  show  many 
peculiarities  due  to  their  being  very  little  ionized.  Salts  as  a  class 
are  highly  ionized  bodies,  and  those  of  mercury  and,  to  a  less  degree, 
those  of  cadmium  are  the  only  conspicuous  exceptions. 

The  mercury  salts  of  volatile  acids,  like  the  corresponding  salts  of 
ammonium  (p.  520),  can  all  be  volatilized  completely.  Mercury 
vapor  and  mercury  compounds  are  poisonous,  the  soluble  ones  more 
markedly  so  than  the  insoluble  ones. 

The  mercurous  salts,  as  a  rule,  are  formed  when  excess  of  mercury 
is  employed,  and  mercuric  salts  when  excess  of  the  oxidizing  acid  or 
other  substance  is  present.  Reducing  agents  turn  mercuric  into  mer- 
curous salts,  and  oxidizing  agents  do  just  the  reverse. 

As  in  the  case  of  the  cuprous  compounds,  it  is  a  question  whether 
simple  or  multiple  formulae,  HgCl  or  Hg2Cl2,  etc.,  should  be  employed 
for  mercurous  salts.  Pending  the  discovery  of  some  basis  for  a  de- 
cision, the  simple  formulae  are  used  here. 

Occurrence  and  Isolation  of  the  Metal.  —  Mercury  occurs 
native  and  as  red,  crystalline  cinnabar,  mercuric  sulphide  HgS.  The 
production  (1913)  in  iron  flasks  (75  Ibs.  each)  is,  Spain  43,800,  Italy 
29,500,  Austria  26,720,  United  States  20,200,  Mexico  4,410. 

The  liberation  of  the  metal  is  easy,  because,  when  roasted,  the 
sulphide  is  decomposed,  and  the  sulphur  forms  sulphur  dioxide.  The 


776  INORGANIC   CHEMISTRY 

mercury  does  not  unite  with  the  oxygen,  for  the  oxide  decomposes 
(p.  17)  at  400-600°:  HgS  +  02  ->  Hg  +  SO2.  In  some  places  the 
ore  is  spread  on  perforated  brick  shelves  in  a  vertical  furnace,  and 
the  gases  pass  through  tortuous  flues  in  which  the  vapor  of  the  metal 
condenses.  The  product  is  filtered  through  chamois-skin.  For  sepa- 
ration from  metallic  impurities,  like  zinc,  arsenic,  and  tin,  which  are 
dissolved,  it  must  be  distilled.  In  the  laboratory,  where  mercury 
finds  many  applications,  it  becomes  impure  with  use,  and  then  adheres 
to  glass,  and  does  not  run  freely  in  spherical  droplets.  For  purifica- 
tion it  is  placed  along  with  a  little  diluted  nitric  acid  in  a  separatory 
funnel  (Fig.  74,  p.  180),  and  kept  in  continual  agitation  by  means  of 
a  current  of  air  drawn  or  blown  through  the  mass,  By  this  treatment, 
foreign  metals,  such  as  sodium  or  zinc,  nearly  all  of  which  are  much 
more  active  than  mercury  (cf.  p.  404),  are  converted  into  nitrates. 
Pure,  dry  mercury  can  be  drawn  off,  when  needed,  at  the  bottom. 
If  a  high  degree  of  purity  is  required,  the  product  must  be  distilled 
in  vacuo. 

Physical  Properties.  —  Mercury  or  quicksilver  (N.L.  hydrargy- 
rum, from  Gk.  v8o>p?  water,  and  apyv/oos,  silver)  is  a  silver-white 
liquid.  At  —38.9°  it  freezes,  and  at  357°  it  boils.  At  ordinary  tem- 
peratures it  has  a  measurable  vapor  tension,  at  15°  0.0008  mm.,  and 
at  100°  0.28  mm.  The  vapor  is  colorless,  does  not  conduct  electricity, 
and  is  monatomic.  A  gold-leaf  suspended  over  mercury  becomes 
amalgamated,  since  the  solution  of  gold  in  mercury  has  a  vapor  ten- 
sion smaller  than  that  of  pure  mercury  (p.  197). 

On  account  of  its  high  density  (13.6,  at  0°)  and  low  vapor  tension, 
the  metal  is  employed  for  filling  barometers  and  manometers.  Its 
uniform  expansion  favors  its  use  in  thermometers.  The  tendency  to 
form  amalgams,  which  it  exhibits  towards  all  the  familiar  metals  with 
the  exception  of  iron  and  platinum  (both,  however,  are  "wet"  by 
it) ,  is  taken  advantage  of  in  various  ways.  Sodium  amalgam  (p.  684) , 
which  is  solid  when  the  sodium  exceeds  2  per  cent,  and  consists  mainly 
of  NaHg2,  behaves  like  free  sodium,  but  with  moderated  activity.  A 
layer  of  mercury  on  the  zinc  plates  of  batteries  reduces  the  action  of 
the  acid  on  the  zinc,  while  the  cells  are  not  in  use.  Mixtures  of  mer- 
cury with  powdered  tin,  silver,  and  gold  quickly  form  solid  amalgams, 
and  are  used  by  dentists.  The  employment  of  mercury  in  the  ex- 
traction of  gold  has  been  mentioned  (p.  758.  See  also  p.  750). 
Mirrors  backed  with  a  tin-mercury  amalgam  have  been  displaced 
by  silvered  mirrors  (p.  755). 


GLUCINUM,   MAGNESIUM,   ZINC,   CADMIUM,   MERCURY     777 

Chemical  Properties.  -  —  When  kept  at  a  temperature  near  to 
its  boiling-point,  mercury  combines  slowly  with  oxygen.  Its  inac- 
tivity towards  oxygen  when  cold  places  it  next  to  the  noble  metals. 
On  account  of  its  general  inactivity,  it  is  used  in  the  laboratory  for 
confining  gases.  It  does  interact  with  hydrogen  sulphide  and  hydro- 
gen iodide,  however  (cf.  Silver,  p.  751).  Mercury  does  not  displace 
hydrogen  from  dilute  acids  (p.  404),  but  with  oxidizing  acids  like 
nitric  acid  and  hot  concentrated  sulphuric  acid,  the  nitrates  and  sul- 
phates are  formed.  With  excess  of  mercury,  the  mercurous  salts,  and 
with  excess  of  the  hot  acid,  the  mercuric  salts,  are  produced.  When 
triturated,  so  that  the  mercury  is  divided  into  minute  droplets  with 
relatively  large  surface,  it  is  used  in  medicine  (blue  pills),  and  shows 
an  activity  which  is  entirely  wanting  in  larger  masses. 

Mercurous  Chloride  HgCL  —  This  salt  (calomel)  is  obtained 
as  a  white  powder  by  precipitation  from  solutions  of  mercurous  salts. 
It  is  manufactured  by  subliming  mercuric  chloride  with  mercury: 


or  more  usually  by  subliming  a  mixture  of  mercuric  sulphate,  made  as 
described  above,  with  mercury  and  common  salt.  It  is  deposited  on 
the  cool  part  of  the  vessel  as  a  fibrous  crystalline  mass.  It  is  slowly 
affected  by  light  just  as  is  silver  chloride.  Here,  however,  the 
chlorine  which  is  released  combines  with  another  molecule  of  the  salt 
to  form  mercuric  chloride.  Since  the  vapor  pressure  of  calomel 
reaches  760  mm.  before  the  temperature  has  risen  to  the  melting-point, 
the  compound  sublimes  at  atmospheric  pressure  without  melting. 
Its  vapor  density  corresponds  to  the  formula  HgCl,  but  the  vapor 
was  shown  by  Smith  and  Menzies  to  consist  wholly  of  Hg  +  HgCl2. 
The  completeness  of  the  dissociation  was  ascertained  by  measuring 
the  partial  pressures  of  the  mercury  and  mercuric  chloride  in  the 
vapor,  and  finding  that  together  they  equalled  the  total  dissociation 
pressure.  The  action  is  reversed  when  the  temperature  falls  (cf. 
Ammonium  chloride,  p.  720).  The  substance  is  used  in  medicine  on 
account  of  its  tendency  to  stimulate  all  organs  producing  secretions. 


Mercuric  Chloride  HgClz.  —  By  direct  union  with  chlorine  the 
mercuric  salt,  corrosive  sublimate  HgCl2,  is  formed.  It  is  usually 
manufactured  by  subliming  mercuric  sulphate  with  common  salt, 
and  crystallizes  in  white,  rhombic  prisms.  It  melts  at  265°  and  boils 
at  307°.  The  solubility  at  20°  is  7.4  :  100  Aq,  and  at  100°,  54  :  100 


778  INORGANIC  CHEMISTRY 

Aq.  It  is  more  soluble  in  alcohol  and  in  ether.  The  aqueous  solu- 
tion is  slightly  acid  in  reaction.  The  salt  is  easily  reduced  to  mer- 
curous  chloride.  When  excess  of  stannous  chloride  is  added  to  the 
solution,  the  white  precipitate  of  calomel,  first  formed,  passes  into 
a  heavy  gray  precipitate  of  finely  divided  mercury  : 

2HgCl2  +  SnCl2  ->  SnCl4  +  2HgCl, 
2HgCl  +  SnCl2  ->  SnCU  +  2Hg. 


The  halides  of  mercury  are  very  little  ionized  in  solution,  the 
bromide  and  iodide  even  less  so  than  the  chloride.  Hence  these  salts 
are  little  affected  by  sulphuric  acid  or  nitric  acid.  For  example,  the 
chlorine  and  nitrosyl  chloride  which  hydrochloric  acid  forms  with  the 
nitric  acid  (p.  537)  are  not  observed  when  mercuric  chloride  is  added 
to  nitric  acid.  On  this  account,  too,  the  hydrolysis  of  the  chloride  is 
much  less  than  that  of  the  nitrate.  There  is  a  tendency  also  to  the 
formation  of  complex  salts,  so  that  the  addition  of  sodium  chloride 
increases  the  solubility  in  water  and  renders  the  solution  neutral, 
NaHgCls  being  formed.  The  complex  salts,  like  K.HgCl3,H2O, 
H2.HgCl4,7H2O,  and  NH4.HgCl3,H2O,  are  easily  made  by  crystalliza- 
tion from  solution.  The  anions  are  relatively  highly  ionized,  how- 
ever, and  the  behavior  is  intermediate  between  that  of  complex  salts 
and  double  salts  (p.  649). 

Corrosive  sublimate,  when  taken  internally,  causes  death.  The 
variable  meaning  of  the  word  "poisonous"  is  well  illustrated  in  this 
case.  The  mercuric  chloride  does  not  act  as  a  direct  poison.  It 
causes  changes  in  the  cells  of  the  kidneys  so  that,  after  about  two 
weeks,  the  waste  products  from  the  system  (urea,  etc.)  can  no  longer 
be  eliminated,  and  death  occurs  from  a  sort  of  autointoxication.  A 
very  dilute  solution  is  used  in  surgery  to  destroy  lower  organisms  and 
thus  prevent  infection  of  wounds.  The  pharmaceutical  tabloids  of 
mercuric  chloride  contain  sodium  chloride,  because,  although  the 
latter  diminishes  the  activity  of  the  compound,  it  also  does  away 
with  the  formation  of  insoluble,  basic  chlorides  and  hastens  solution. 
Mercuric  chloride  acts  also  as  a  preservative  of  zoological  materials, 
forming  insoluble  compounds  with  proteins,  and  preventing  their 
decay.  For  the  same  reason,  albumin  (white  of  an  egg,  a  mixture  of 
proteins)  is  given  as  an  antidote  in  cases  of  sublimate  poisoning. 

The  Iodides  of  Mercury.  —  Mercurous  iodide  Hgl  is  formed 
by  rubbing  iodine  with  excess  of  mercury.  It  also  appears  as  a  green- 


GLUCINUM,  MAGNESIUM,  ZINC,  CADMIUM,   MERCURY     779 

ish-yellow  precipitate  when  potassium  iodide  is  added  to  a  solution 
of  a  mercurous  salt.  The  compound  decomposes  spontaneously  into 
mercury  and  mercuric  iodide.  The  decomposition  is  much  hastened 
by  the  use  of  excess  of  potassium  iodide,  which  combines  with  and 
removes  the  mercuric  iodide  (see  below)  : 


Mercuric  iodide  HgI2  is  obtained  by  direct  union  of  mercury  with 
excess  of  iodine,  or  by  addition  of  potassium  iodide  to  a  solution  of  a 
mercuric  salt.  It  is  a  scarlet  powder,  insoluble  in  water,  but  soluble 
in  alcohol  and  ether.  It  interacts  with  excess  of  potassium  iodide, 
forming  the  soluble,  colorless  potassium  mercuri-iodide  K2.HgI4  with 
which  many  precipitants  fail  to  give  mercury  compounds.  When 
heated  above  128°  it  turns  into  a  yellow  modification,  and  at  223° 
this  new  form  melts.  On  being  cooled,  the  liquid  freezes  first  in  the 
tetragonal  yellow  form,  and  below  128°,  especially  if  touched  with  a 
glass  rod,  it  changes  into  the  red,  monoclinic  variety  with  evolution  of 
heat.  Sulphur  (p.  411)  and  ammonium  nitrate  (p.  679)  show  similar 
transition  points.  When  the  vapor  of  the  compound  is  cooled,  it 
first  forms  thin  scales  of  the  yellow  form,  which  is  the  unstable  one  at 
low  temperatures,  and  these  turn  red  when  touched.  Similarly,  pre- 
cipitation gives  first  the  yellow  variety,  which  presently  becomes  red 
(cf.  Transformation  by  steps,  p.  544). 

The  Oxides.  —  When  bases  (excepting  ammonium  hydroxide. 
see  p.  781)  are  added  to  solutions  of  mercurous  salts,  the  brownish- 
black  mercurous  oxide  Hg2O  is  thrown  down.  The  hydroxide  is 
doubtless  formed  transitorily  and  then  loses  water  (cf.  Silver  oxide, 
p.  753).  Under  the  influence  of  light  or  gentle  heat  (100°),  this 
compound  resolves  itself  into  mercuric  oxide  and  mercury. 

Mercuric  oxide  HgO  is  formed  as  a  red,  crystalline  powder,  when 
mercury  is  heated  in  air  near  to  357°,  but  is  usually  made  by  decom- 
posing the  nitrate.  Commercial  specimens,  incompletely  decom- 
posed, thus  frequently  give  the  brown  nitrogen  tetroxide  when  heated. 
It  is  formed  also  as  a  yellow  powder  by  adding  bases  (excepting 
ammonium  hydroxide,  see  p.  781)  to  solutions  of  mercuric  salts.  It 
is  contended  by  some  chemists  that  the  difference  in  activity  between 
the  red  and  yellow  forms  is  due  solely  to  the  finer  state  of  division 
of  the  latter,  and  by  others  it  is  maintained  that  the  substance  is 
dimorphous  (p.  412)  and  that  two  distinct  varieties  exist. 


780  INORGANIC  CHEMISTRY 

The  Nitrates.  —  The  mercurous  salt  HgN03,H20  is  formed  by 
the  action  of  cold,  diluted  nitric  acid  upon  excess  of  mercury.  It 
forms  monoclinic  crystals.  It  is  hydrolyzed,  slowly  by  cold,  and 
rapidly  by  warm  water,  giving  an  insoluble  basic  nitrate: 

2HgN03  +  H20  <=±  HN03  +  Hg2(OH)NO3  J, . 

On  this  account  a  clear  solution  can  be  made  only  when  some  nitric 
acid  is  added.  Free  mercury  is  also  kept  in  the  solution  to  reduce 
mercuric  nitrate,  which  is  formed  by  atmospheric  oxidation: 

Hg(N03)2  +  Hg  ->  2HgN03    or    Hg++ +  Hg ->  2Hg+. 

Mercuric  nitrate  Hg(NO3)2,8H20  is  produced  by  using  excess  of 
warm,  concentrated  nitric  acid.  It  forms  rhombic  tables.  The 
aqueous  solution  is  strongly  acid,  and  deposits  a  yellowish,  crystal- 
line, basic  nitrate  Hg3(OH)2O(NO3)2.  The  hydrolysis  is  reversed  by 
adding  nitric  acid. 

Sulphides  of  Mercury.  —  Mercurous  sulphide  Hg2S  is  formed 
by  precipitation  from  mercurous  salts,  but  is  stable  only  below  — 10°. 
Above  — 10°  it  decomposes  into  mercury  and  mercuric  sulphide. 

Crystallized  mercuric  sulphide  HgS  occurs  as  cinnabar,  and  is  red. 
When  formed  by  precipitation  with  hydrogen  sulphide,  or  by  rubbing 
together  mercury  and  sulphur,  it  is  black  and  amorphous.  By  subli- 
mation, in  the  course  of  which  it  dissociates,  the  black  form  gives  the 
red,  crystalline  one.  When  allowed  to  stand  under  a  solution  of  so- 
dium sulphide,  the  black  form  is  slowly  transformed  into  the  red. 
This  shows  that  the  red  form  is  the  more  stable,  possesses  less  energy, 
and  is  less  soluble  at  ordinary  temperatures.  The  change  is  effected 
by  intermediate  formation  of  a  complex  sulphide,  the  solution,  when 
saturated  toward  the  less  stable  black  sulphide,  being  supersaturated 
toward  the  more  stable  red  one.  A  white,  crystalline  sodium  mercuri- 
sulphide  Na2HgS2,8H20  can  be  obtained  from  the  solution. 

The  black  and  the  red  varieties  .do  not  interact  even  with  boiling 
nitric  acid,  which  oxidizes  most  sulphides  readily.  They  are,  there- 
fore, less  soluble  in  water  than  is  cupric  sulphide  (pp.  421,  774). 
They  are  attacked,  however,  by  aqua  regia,  because  of  the  formation 
of  the  negative  ion  (see  gold,  p.  759)  of  a  complex  salt  H2.HgCl4 
(  =  2HCl,HgCl2). 

The  red  form  of  the  sulphide  is  used  in  making  paint  (vermilion). 
The  color  is  more  permanent  than  that  of  red  lead  Pb3O4,  because  re- 
ducing gases  (e.g.,  SCfe),  acids  (e.g.,  H2SO4),  and  hydrogen  sulphide, 


GLUCINUM,   MAGNESIUM,   ZINC,   CADMIUM,   MERCURY     781 

which  are  present  in  the  air,  do  not  affect  it.  It  is  not  stable,  how- 
ever, when  applied  to  metals,  since  iron,  zinc,  etc.,  all  displace  mer- 
cury from  combination,  and  in  these  cases  red  lead  is  preferred. 

Mercuric  Cyanide  Hg(NC)z.  —  This  salt  is  made  by  treating 
precipitated  mercuric  oxide  with  hydrocyanic  acid,  and  is  obtained  in 
square-prismatic  crystals.  When  heated  it  gives  off  cyanogen: 
Hg(NC)2  — >  Hg  -}-  C2N2,  and  is  a  convenient  source  of  this  gas  (p.  625) . 
The  compound  is  soluble  in  alcohol,  ether,  and  water.  In  solution  in 
water  it  is  so  little  ionized  that  the  freezing-point  of  the  solution  is 
normal  (p.  336) ,  and  many  reagents  fail  to  show  the  presence  of  either 
ion.  Thus,  with  silver  nitrate  no  silver  cyanide  is  precipitated,  and 
with  a  base  no  mercuric  oxide.  With  potassium  cyanide  it  forms  a 
complex  cyanide  K2.Hg(CN)4.  Hydrogen  sulphide  throws  down  the 
sulphide  from  both  the  simple  and  the  complex  cyanides. 

The  Fulminate  and  Thiocyanate.  —  Mercuric  fulminate 
Hg(ONC)2  is  obtained  as  a  white  precipitate  when  mercury  is  treated 
with  nitric  acid,  and  alcohol  is  added  to  the  solution.  It  decomposes 
suddenly  when  struck,  and  is  used  in  making  percussion  caps. 

The  thiocyanate  Hg(NCS)2  is  precipitated  when  potassium  thio- 
cyanate  K(NCS)  is  added  to  a  solution  of  mercuric  nitrate.  When 
formed  into  little  balls  and  burned  in  the  air,  the  substance  leaves  a 
curiously  voluminous  ash  ("Pharaoh's  serpents"). 

Ammono-  Compounds  of  Mercury.  —  When  ammonium  hy- 
droxide is  added  to  a  solution  of  a  mercuric  salt,  a  white  substance, 
of  a  type  which  we  have  not  previously  encountered,  is  thrown  down. 
Mercuric  chloride  gives  Hg(NH2)Cl,  commonly  called  "infusible 
white  precipitate,"  or  ammono-basic  mercuric  chloride. 

HgCl2  +  H.NH2  +  NH3  ->  Hg(NH2)Cl  +  NH4C1. 

The  action  is  similar  to  an  hydrolysis  which  gives  a  basic  salt: 
HgCl2  +  H.OH  ->  Hg(OH)Cl  +  HC1,  excepting  that  ammonia 
H.NH2  plays  the  part  of  the  water.  Water  gives  aquo-basic  salts. 
When  liquid  ammonia  is  the  solvent,  ammono-basic  salts  are  pro- 
duced. In  a  few  cases,  as  here,  an  ammono-basic  salt  is  obtained 
even  when  water  is  present.  The  study  of  reactions  in  liquid  am- 
monia solutions  by  E.  C.  Franklin  has  led  to  the  discovery  of  a  large 
number  of  new  and  most  interesting  substances. 

The  addition  of  ammonium  hydroxide  to  a  solution  of  potassium 


782  INORGANIC   CHEMISTRY 

mercuri-iodide  K2HgI4  gives  rise  to  a  compound  of  the  same  type, 
ammono-basic  mercuric  iodide  Hg2NI,H20,  which  appears  as  a 
brown  precipitate: 

2HgI2  +  H3.N  +  3NH3  ->  Hg2NI  +  3NH4I. 

A  solution  of  potassium  mercuri-iodide  containing  potassium  hydrox- 
ide, Nessler's  reagent,  becomes  distinctly  yellow  with  traces  of 
ammonia,  and  brown  with  larger  amounts,  and  is,  therefore,  a  valu- 
able reagent  for  detecting  traces  of  this  base. 

Mercuric  nitrate  Hg(N03)2  and  ammonium  hydroxide  give  an 
insoluble  ammono-basic  mercuric  nitrate,  Hg  =  N  — HgNO3  which  is 
more  basic  than  the  ammono-basic  chloride: 

2Hg(N03)2  +  H3.N  +  3NH3  ->  Hg2(N)NO3  +  3NH4NO3. 

When  calomel  is  treated  with  ammonium  hydroxide,  it  turns  into 
a  black,  insoluble  body.  This  is  a  mixture  of  free  mercury,  to 
which  it  owes  its  dark  color,  and  " infusible  white  precipitate," 
Hg  +  Hg(NH2)Cl.  To  this  reaction  calomel  owes  its  name  (Gk. 
KaA.o/x,e'Aas,  beautiful  black).  Mercurous  nitrate  gives  a  black,  in- 
soluble mixture,  2Hg  +  Hg2(N)N03. 

Calomel,  when  dry,  absorbs  ammonia  gas,  forming  a  molecular 
compound  of  the  common  type  HgCl,NH3.  This  substance  loses 
the  ammonia  again  when  the  pressure  is  reduced.  The  other  com- 
pounds described  above,  on  the  other  hand,  do  not  contain  nitrogen 
and  hydrogen  in  the  proportions  necessary  to  form  ammonia  and  are 
stable.  Hence  they  are  necessarily  to  be  regarded  as  belonging  to  a 
different  type. 

Analytical  Reactions  of  Mercury  Compounds.  —  The  two 

ionic  forms  of  the  element,  mercurous-ion  Hg+  and  mercuric-ion 
Kg"1"*",  are  both  colorless.  Their  chemical  behavior  is  entirely 
different.  Both  give  the  black  sulphide  HgS,  which  does  not  inter- 
act with  acids  and  other  solvents  of  mercury  salts.  Mercurous-ion 
gives  the  insoluble,  white  chloride,  the  black  oxide,  and  a  black 
mixture  with  ammonium  hydroxide.  Mercuric-ion  gives  a  soluble 
chloride,  a  yellow,  insoluble  oxide,  and  a  white  precipitate  with  am- 
monium hydroxide.  The  behavior  with  stannous  chloride  (p.  778)  is 
characteristic.  With  potassium  iodide  the  two  ions  behave  differ- 
ently (p.  778).  The  more  active  metals  displace  mercury  from  all 
compounds.  Copper  is  used  as  the  displacing  metal,  in  testing  for 
Hg+  or  Hg++,  because  the  silvery  mercury  is  easily  seen  on  its  surface. 


GLUCINUM,   MAGNESIUM,   ZINC,   CADMIUM,   MERCURY     783 

Salts  of  mercury  are  volatile.  When  heated  in  a  tube  with 
sodium  carbonate,  they  give  a  sublimate  of  metallic  mercury. 

THE  RECOGNITION  OF  CATIONS  IN  QUALITATIVE  ANALYSIS 

" Wet-way"  analysis  consists  in  recognizing  the  various  positive 
and  negative  ions  present  in  a  solution  (p.  385).  It  was  stated  that 
the  sulphides  might  be  divided  into  three  classes  according  to  their 
behavior  towards  water  and  acids  (pp.  421,  774).  Now  these  differ- 
ences in  behavior  furnish  us  with  a  basis  for  distinguishing  the  cations 
present  in  a  solution.  Since  the  properties  of  many  sulphides  and 
other  compounds  of  the  metals  have  been  studied  in  recent  chapters, 
it  is  now  possible  to  make  a  more  complete  statement. 

The  following  plan,  taken  in  conjunction  with  the  statements  in 
the  context,  shows  how  a  single  cation  may  be  identified,  and  how, 
when  several  cations  are  present,  a  separation  preparatory  to  identifi- 
cation may  be  effected.  What  will  be  said  applies  only  to  the  case  of 
a  solution  containing  salts  like  the  chlorides,  nitrates,  or  sulphates  of 
one  or  more  cations,  and  leaves  the  oxalates,  phosphates,  cyanides, 
and  some  other  salts,  out  of  consideration. 

Group  1.  —  It  is  usual  to  add,  first,  hydrochloric  acid,  to  find  out 
whether  cations  giving  insoluble  chlorides  are  present.  Argentic, 
mercurous,  and  plumbic  salts  give  the  white  AgCl,  HgCl,  and  PbCl2, 
respectively.  (For  the  further  recognition  of  each,  see  p.  757.) 
Filtration  eliminates  the  precipitate,  if  there  is  any. 

Group  2.  —  A  free,  active  acid  being  now  present,  hydrogen  sul- 
phide is  led  into  the  solution.  The  sulphides  insoluble  in  active  acids, 
namely,  HgS,  CuS,  PbS,  Bi2S3,  CdS,  As2S3,  Sb2S3,  SnS,  are  therefore 
thrown  down.  The  first  four  are  black  or  brown,  the  next  two  are 
yellow,  and  the  last  two  are  orange  and  brown  respectively.  A  dark- 
colored  substance  will  naturally  obscure  one  of  lighter  color,  if  more 
than  one  is  present.  If  too  much  acid  is  used,  the  precipitation  of 
several  of  the  sulphides  will  be  incomplete  (p.  774) ;  if  too  little,  zinc 
sulphide  may  come  down  (p.  772).  Filtration  again  eliminates  the 
precipitate. 

This  group  is  easily  subdivided.  Any  or  all  of  the  last  three  sul- 
phides will  pass  into  solution  when  warmed  with  yellow  ammonium 
sulphide,  for  they  give  soluble  complex  sulphides  similar  to  potassium 
sulphaurate  (p.  760).  The  first  five  sulphides,  or  any  of  them,  will  be 
unaffected.  On  the  other  hand,  these  five  sulphides,  with  the  excep- 
tion of  HgS,  will  interact  with  hot  nitric  acid  (p.  774).  Other  reac- 


784  INORGANIC   CHEMISTRY 

tions  described  in  the  context  are  then  used  to  distinguish  between,  or, 
if  there  is  a  mixture,  to  separate,  the  members  of  the  sub-groups. 

Group  3.  —  The  solution  (filtrate)  is  now  neutralized  with  am- 
monium hydroxide,  and  ammonium  sulphide  is  added.  Some  am- 
monium chloride  is  also  used,  to  prevent  the  precipitation  of  magne- 
sium hydroxide  (p.  766),  which,  in  any  event,  would  be  incomplete. 
The  sulphides  which  are  insoluble  in  water,  and  are  not  hydrolyzed  by 
it,  now  appear.  They  are  FeS,  CoS,  NiS,  all  black,  MnS  and  ZnS, 
which  are  pink  and  white,  respectively.  There  are  precipitated  also 
the  hydroxides  of  chromium  and  of  aluminium,  Cr(OH)3  and  A1(OH)3, 
because  their  sulphides  are  hydrolyzed  by  water,  and  the  hydroxides 
are  formed  by  the  hydroxide-ion  in  the  ammonium  sulphide  solution 
(cf.  p.  648).  They  are  too  insoluble  to  behave  like  magnesium  hy- 
droxide (p.  766)  by  dissolving  in  salts  of  ammonium.  They  also 
form  no  complex  metal-ammonia  ions,  as  does  zinc  (p.  771).  The 
sulphides  of  nickel  and  cobalt  resemble  the  sulphide  of  zinc  »in  being 
precipitated  by  hydrogen  sulphide  when  acetic  acid  is  the  only 
acid  present.  The  other  sulphides  interact  even  with  acetic  acid 
(p.  774). 

Another  plan  is  to  oxidize  the  iron,  if  present,  and  use  ammonium 
chloride  and  ammonium  hydroxide  instead  of  ammonium  sulphide. 
The  hydroxides  of  the  trivalent  elements,  Fe(OH)3,  Cr(OH)8,  A1(OH)8, 
can  be  precipitated  by  excess  of  ammonium  hydroxide,  even  when 
salts  of  ammonium  are  present.  Those  of  the  bivalent  metals, 
Mn(OH)2,  Fe(OH)2,  Zn(OH)2,  Ni(OH)2,  Co(OH)2,  resemble  magne- 
sium hydroxide  (p.  766),  and,  of  these,  the  last  three  resemble  also 
zinc  hydroxide  (p.  771),  and  so  cannot  be  precipitated.  After  ni- 
tration, ammonium  sulphide  now  throws  down  the  sulphides  of  the 
five  bivalent  metals  (for  a  third  plan,  see  Chemical  relations  of 
aluminium) . 

Group  4.  —  After  filtration  from  members  of  the  iron  group,  if 
any  were  present,  ammonium  carbonate  is  added,  and  precipitates  the 
remaining  metals  whose  carbonates  are  insoluble,  BaCO3,  SrCO3, 
CaCO3,  with  the  exception  of  magnesium  (p.  767). 

By  addition  of  ammonium  phosphate  to  a  portion  of  the  filtrate, 
magnesium,  if  present,  now  comes  out  in  the  form  NH4MgPO4. 
There  remain  in  solution  only  salts  of  potassium,  sodium,  and  am- 
monium. Since  only  ammonium  compounds,  and  other  sub- 
stances which  can  be  volatilized  have  been  added,  evaporation  and 
ignition  of  the  residue  leaves  the  salts  of  the  two  metals.  If  no  other 
metallic  elements  have  been  shown  to  be  present,  it  saves  time  to 


GLUC1NUM,   MAGNESIUM,  ZINC,   CADMIUM,   MERCURY     785 

examine  a  fresh  portion  of  the  original  material.     Salts  of  ammonium 
must  also  be  sought  in  a  fresh  sample  by  the  usual  test  (p.  681). 

The  following  simple  compounds  are  soluble,  but  are  so  little 
ionized  that  their  solutions  do  not  show  all  the  reactions  of  both  of 
the  ions:  NH4OH,  H2S,  HNC,  H2CO3,  HgCl2,  Hg(CN)2,  Fe(NCS)3, 
Pb(C2H3O2)2.  With  a  number  of  others,  for  example  CdI2,  the  actions 
are  incomplete  for  the  same  reason.  Complex  compounds,  as  we 
have  seen,  give  complex  ions,  and  these  ions  are  usually  so  little  re- 
solved into  simpler  ions  that  the  latter  cannot  be  discovered  by  all 
the  usual  tests.  Thus:  K.Ag(CN)2  gives  K+  and  Ag(CN)2-,  but 
very  little  Ag+  and  CN~  (p.  752);  Cu(NH3)2.Cl2  gives  much 
Cu(NH3)2++,  and  Cl~  but  very  little  Cu++.  The  individual  cases 
are  described  in  the  context. 

Exercises.  —  1.  What  is  the  numerical  value,  (a)  of  the  solu- 
bility product  of  magnesium  hydroxide,  (6)  of  the  concentration  of 
hydroxide-ion  given  by  it  and  by  normal  ammonium  hydroxide, 
respectively  (Table)?  Will  normal  concentration  of  ammonium 
chloride  suffice  to  reduce  the  latter  below  the  former? 

2.  Why  should  we,  perhaps,  expect  ammonium  sulphide  solution 
to  precipitate  magnesium  hydroxide,  and  why  does  it  not  do  so? 

3.  What  volume  of  air  is  required  to  oxidize  one  formula-weight 
of  zinc  sulphide  to  ZnO  and  SO2,  and  what  volume  of  sulphur  dioxide 
is  produced?     Is  the  product  more  or  less  diluted  with  nitrogen  than 
when  pure  sulphur  is  burned,  and  by  how  much? 

4.  Make  equations  showing,  (a)  the  effect  of  heating  zinc  chloride 
with  cobalt  nitrate  Co(NO3)2  in  the  Bunsen  flame  (p.  771),  (6)  the 
action  of  hydrogen  sulphide  on  sodium  zincate,  (c)  the  actions  of  con- 
centrated nitric  acid  and  of  concentrated  sulphuric  acid  on  mercury. 

5.  What  is  the  distinction  between  a  solid  isomorphous  mixture  of 
two  salts  and  a  double  salt? 

6.  What  kind  of  salts  might  take  the  place  of  sodium  acetate  in 
the  precipitation  of  zinc  sulphide  (p.  772)?     Give  examples. 

7.  Compare  the  amalgamation  of  a  gold-leaf  by  mercury  vapor 
with  the  phenomenon  of  deliquescence  (p.  776). 

8.  If  the  scheme  for  the  recognition  of  cations  (p.  783)  were  ap- 
plied to  solutions  prepared  from  materials  containing,  (a)  calcium  oxa- 
late  and,  (6)  potassium  argenticyanide,  at  what  stage  and  how  would 
the  presence  of  each  of  these  substances  affect  the  normal  order? 

9.  Why  do  none  of  the  salts  of  the  elements  in  this  family  give 
recognizable  effects  with  the  borax  bead? 


CHAPTER  XXXVIII 
ELECTROMOTIVE   CHEMISTRY 

WE  have  seen  that  many  chemical  changes  are  accompanied  by 
a  liberation  of  energy.  If  no  special  arrangement  is  made,  the  energy 
is  always  liberated  in  the  form  of  heat,  light,  and  mechanical  energy. 
In  changes  involving  ionogens,  however,  the  energy  can  be  secured 
in  the  form  of  electricity.  Most  changes  between  ionogens,  including 
oxidations,  may  be  adapted  so  as  to  deliver  this  form  of  energy.  It 
need  hardly  be  added  that,  since  the  transformation  of  chemically 
equivalent  amounts  of  different  sets  of  substances  produces  very  dif- 
ferent quantities  of  heat,  so  it  produces  also  correspondingly  different 
amounts  of  electrical  energy.  Thus  the  original  free  internal  energy 
may,  theoretically,  be  measured  by  either  method.  In  practice, 
however,  the  thermochemical  plan  fails  entirely  in  many  cases  (cf. 
pp.  35,  691),  and  the  electrical  is,  as  we  shall  see,  often  much  more 
instructive.  The  study  of  what,  to  parody  the  phraseology  of 
thermochemistry,  we  might  call  exelectrical  actions,  thus  resolves 
itself  into  constructing  experimental  battery-cells  involving  all  kinds 
of  chemical  changes,  and  studying  the  electric  currents  which  are  set 
in  motion  by  the  progress  of  the  changes.  We  have  therefore  named 
this  branch  of  the  science  electromotive  chemistry. 

In  addition  to  its  significance  theoretically,  for  the  purpose  of 
measuring  chemical  affinity  in  ionic  actions,  electromotive  chem- 
istry has  recently  acquired  great  commercial  importance  because  of 
the  rapid  multiplication  of  electro-chemical  industries.  It  is  true 
that  the  majority  of  the  actions  used  in  these  industries  are  electro- 
lytic (endelectrical),  and  that  this  sort  of  change  is  the  precise  inverse 
of  the  other,  since  in  it  electricity  is  consumed  instead  of  being  set 
in  motion,  but  it  is  also  true  that  neither  variety  can  be  understood 
without  a  study  of  both. 

Factors  and  Units  of  Electrical  Energy.  —  On  account  of  the 
close  relation  between  electromotive  chemistry  and  electrolysis,  parts 
of  the  former  subject  were  anticipated  when  the  latter  was  discussed 
(pp.  357-361).  These  pages  should  now  be  re-read  attentively.  In 

786 


ELECTROMOTIVE  CHEMISTRY  787 

particular,  it  must  be  recalled  that  a  quantity  of  electrical  energy  is 
expressed  by  two  factors.  One  is  called  the  .quantity  of  electricity, 
and  is  measured  in  coulombs.  The  other  is  called  the  electromotive 
force  in  the  case  of  a  current,  or,  when  a  current  is  not  flowing  or  is 
not  being  considered,  the  difference  in  potential,  and  is  expressed  in 
volts.  Just  as  in  electrolysis  chemically  equivalent  quantities  of  ele- 
ments or  ions,  in  being  liberated  from  solutions  of  different  substances, 
use  up  equal  quantities  of  electricity  (Faraday's  law,  p.  350),  so  in  a 
battery-cell  the  interaction  of  chemically  equivalent  amounts  of  dif- 
ferent sets  of  substances  produces  equal  quantities  of  electricity  (p. 
353).  Likewise,  just  as  in  the  former  case  different  amounts  of  elec- 
trical energy  (p.  359),  and  therefore  different  electromotive  forces, 
are  required  to  produce  in  different  solutions  equivalent  amounts  of 
chemical  change  (p.  216),  so  in  the  latter  case  different  amounts  of 
electrical  energy  are  generated  by  the  complete  interaction  of  chemical 
equivalents  of  different  sets  of  substances,  and  therefore  diverse 
differences  in  potential  are  created  and  currents  of  different  electro- 
motive force  are  produced.  The  electrical  energy  used  in  the  former 
case  or  produced  in  the  latter  is  expressed  by  the  product  of  the 
factors: 

No.  of  coulombs  X  No.  of  volts 

=  Quant,  of  elect,  energy  (in  joules,  p.  34). 

If  we  consider  the  time  occupied  by  either  process,  and  wish  to  express 
the  rate  at  which  the  energy  is  consumed  or  produced,  we  regard  1 
coulomb  per  second  (1  ampere)  as  the  unit.  Hence: 

No.  of  amperes  X  No.  of  volts  =  Joules  per  sec.  =  Watts. 

The  erg  (p.  34)  is  so  small  as  a  unit  of  energy  or  work,  that  the 
joule  (=  10,000,000  ergs)  and  the  kilojoule  (1000  joules)  are  more 
often  employed.  Similarly,  the  rate  at  which  the  energy  is  delivered 
or  used  (the  power)  is  expressed  by  the  watt  (=  10,000,000  ergs  per 
sec.)  or  the  kilowatt  (1000  watts).  The  horsepower  is  746  watts. 

An  illustration  will  show  the  meaning  of  this  relation.  If  a 
50-watt  incandescent  lamp  is  used  on  a  110- volt  circuit,  by  sub- 
stituting these  values  in  the  equation  we  perceive  that  such  a  lamp 
must  carry  about  0.5  amperes,  or  one  coulomb  every  two  seconds. 
If,  with  the  same  voltage,  we  wanted  a  lamp  to  carry  more  electricity 
per  second,  we  should  have  to  reduce  the  resistance  of  the  lamp,  say, 
by  shortening  the  filament,  or  using  a  thicker  one.  Evidently,  the 
number  of  such  lamps  required  to  consume  one  horsepower  would  be 


788  INORGANIC  CHEMISTRY 

746/50,  or  between  14  and  15  lamps.  Again,  to  decompose  one 
molecular  weight  of  hydrochloric  acid  (36.5  g.)  96,504  coulombs 
(p.  357)  are  required,  and  an  E.M.F.  of  at  least  1.83  volts  (see 
p.  798).  The  electrical  energy  needed  is  therefore  96,504  X  1.83  = 
176,602  joules.  If  this  were  to  be  accomplished  by  the  current 
from  a  110-volt  direct-current  lighting  circuit,  passing  through  a 
50-watt  lamp  in  series  with  the  electrolytic  cell,  the  time  required 
(x  seconds)  would  be  given  by  50  joules  per  sec.  X  xsecs.  =  176,602 
joules,  where  x  =  3532  seconds,  or  about  59  minutes. 

The  factors  of  electrical  energy  (volts  and  amperes)  are  easily 
measured  when  electricity  is  produced,  and  are  easily  provided 
according  to  any  specification  when  electricity  is  to  be  used.  Hence, 
it  is  much  easier  to  study  the  relations  between  chemical  change 
and  this  form  of  energy  than  between  the  same  change  and  the 
heat  or  any  other  form  of  energy  which,  under  other  conditions,  it 
might  produce.  Electrochemistry  is,  therefore,  in  many  ways  better 
understood,  and  easier  to  handle  than  are  other  branches  of  chemistry 
involving  energy. 

Some  Reactions  that  can  be  Used  to  Furnish  Electricity. 

—  A  few  illustrations  of  the  kinds  of  reactions  which  can  easily  be 
carried  out  in  cells,  so  as  to  furnish  an  electric  current  instead  of 
heat,  may  be  classified  thus: 

Combination  cells,  such  as  one  in  which  zinc  (or  some  other 
active  metal)  and  bromine  are  the  reacting  substances.  If  zinc 
be  placed  in  bromine-water  (or  with  pure  bromine),  we  obtain 
zinc  bromide: 

Zn  +  Br2  ->  ZnBr2,     or    Zn°  +  2Br°  ->  Zn++  -f  2Br~. 

Displacement  cells,  such  as  one  with  cupric  sulphate  solution 
and  a  metal  more  active  than  copper  (e.g.,  Mg,  Al,  Zn,  or  Fe),  and 
able  to  displace  (p.  403)  this  element: 

Zn  +  CuS04  -»  ZnS04  +  Cu,     or    Zn°  +  Cu++  -»  Zn++  +  Cu°. 
A  non-metal  may  also  be  displaced: 

2KI  +  Br2  ->  2KBr  +  I2,     or    I~  +  Br°  -*  1°  +  Br~. 

Oxidation  cells,  such  as  one  in  which  ferrous  chloride  FeCl2  or 
stannous  chloride  SnCl2  is  oxidized  by  chlorine-water,  giving  FeCl3 
or  SnCl4: 

Cl2-»SnCl4,    or    Sn++  +  2CP  ->  Sn++++  +  2Cr. 


ELECTROMOTIVE   CHEMISTRY 


789 


Concentration  cells,  or  cells  in  which  the  same  substance  in  two 
different  concentrations  is  used. 


The  Arrangement  of  the  Cell.  —  Every  cell  has  one  striking 
characteristic.  If  the  pairs  of  substances  mentioned  in  the  last 
section  are  placed  together,  they  interact  and  heat  is  produced.  There 
is  no  way  to  avoid  the  action,  and  the  liberation  of  the  energy  as 
heat,  if  the  substances  come  in  contact.  If,  therefore,  all  the  energy 
is  to  be  obtained  as  electrical  energy,  the  substances  must  be  pre- 
vented from  coming  in  contact  with  one  another.  Paradoxical  as 
it  may  seem,  it  is  easily  possible  to  obtain  the  electricity,  and  yet 
fulfill  this  essential  condition.  The  plan  in  all  cells  is  to  place  the 
one  substance  in  or  round  one  pole,  and  the  other  substance  in  or 
round  the  other  pole,  and  to  separate  the  substances  by  a  porous 
partition,  or  some  equivalent  arrangement. 

Suppose  that  it  is  the  first  of  the  above-mentioned  actions  that 
is  to  be  used  —  the  action  of  zinc  and  bromine.  The  active  sub- 
stances are  arranged  as  follows:  The  pole 
on  the  left  (Fig.  149)  is  metallic  zinc. 
The  solution  on  the  right  contains  the 
bromine.  The  porous  partition  in  the 
canter  is  permeable  by  migrating  ions, 
but  hinders  the  mere  diffusion  of  the  dis- 
solved bromine  towards  the  zinc,  and  so 
prevents  direct  interaction  with  liberation 
of  heat. 

Now,  to  enable  the  cell  to  operate, 
inactive,  conducting  substances  must  be 
added  to  complete  the  arrangement.  A 
pole  is  added  on  the  right,  a  conducting 
solution  is  placed  to  the  left  of  the  parti- 
tion, and  a  wire  must  connect  the  two 
poles.  The  wire  may  connect  the  poles  through  a  voltmeter  ^  so  that 
the  E.M.F.  produced  may  be  measured.  Also,  since  bromine-water 
is  a  poor  conductor,  a  well-ionized  salt  must  be  present  along  with  the 
bromine.  The  substances  used  for  these  purposes  must  be  inactive. 
For  example,  the  pole  on  the  right  must  be  a  conductor,  but  its 
material  must  not  interact  chemically  with  the  bromine  or  with 
the  salt.  A  rod  of  carbon  or  a  platinum  wire  will  serve  the  purpose. 
A  more  active  metal,  such  as  copper,  could  not  be  used,  because  it 
would  combine  with  the  bromine.  Again,  common  salt  or  sodium 


Pos.  ions  -* 


Neg.  ions^- 


FJG.  149. 


790 


INORGANIC  CHEMISTRY 


nitrate  may  be  mixed  with  the  bromine,  because  it  will  not  interact 
with  bromine  or  carbon  or  platinum.  Still  again,  the  solution  added 
on  the  left  must  be  one  which  will  not  act  upon  the  zinc  pole,  nor 
upon  the  solution  on  the  right,  which  it  meets  inside  the  porous 
partition.  Common  salt  or  niter  fulfills  these  conditions.  An  acid 
could  be  used  on  the  right,  but  not  on  the  left,  for  it  would  interact 
with  the  zinc.  The  reader  should  make  a  different  selection  of  in- 
active materials,  so  as  to  become  familiar  with  the  reasoning  involved 
in  the  choice  in  each  case. 

Note  that  in  each  figure,  the  symbols  for  the  active  substances 
are  in  black-face  type,  the  products  are  in  Roman  type,  and  the 
inactive  materials  are  in  italic  type. 

The  Operation  of  the  Cell.  —  When  the  cell  has  been  assem- 
bled, and  the  wires  have  been  connected,  the  following  phenomena 

are  observed: 

1.  The  zinc  begins  to  form  zinc  ions, 
Zn  — ->  Zn++,  an  operation  which  leaves 
the  pole  negative  (Fig.  150). 

2.  The    bromine    molecules    nearest 
their  pole  touch  this  pole,  become  bro- 
mide ions,  Br2  — »  2Br~,  and  leave  the  pole 
positive. 

3.  Since  one  pole  is  negative  and  the 
other  positive,  a  current  flows  through 
the  wire. 

4.  The  new  positive  ions  (Zn++)  round 
the  left  pole  (anode)  attract  all  the  neg- 
ative ions  in  the  cell,  and  cause  them  to 
migrate  towards  the  left  so  as  to  keep 

all  parts  of  the  solution  electrically  neutral.  They  also  repel  all  the 
positive  ions. 

5.  The  new  negative  ions  on  the  right  (Br~)  similarly  attract  all 
the  positive  ions  in  the  cell,  and  cause  them  to  drift  slowly  towards 
the  right  pole  (cathode).    They  repel  all  the  negative  ions. 

6.  (Very  important.)     It  will  be  seen  that  the  zinc  and  the  bro- 
mine become  ionized  at  a  distance  from  one  another  and  do  not  actu- 
ally combine.     The  slow  migration  of  the  Zn++  and  Br~  ions  will,  of 
course,  after  some  hours  or  days,  bring  some  of  these  ions  together 
in  or  near  the  partition,  and  some  molecules  will  be  formed.     But  this 
operation  produces  no  electrical  energy  —  it  only  gives  out  or  absorbs 


Nad 


ions-*        Neg.  ions-*- 
Fio.  150. 


ELECTROMOTIVE  CHEMISTRY 


791 


heat  (p.  368).  It  is  not  an  essential  part  of  the  operation  of  the  cell. 
The  chemical  change  which  produces  the  current  is  the  ionization 
of  the  two  elements,  separately.  The  term  combination  cell  is, 
therefore,  misleading.  The  cell,  as  a  source  of  electrical  energy,  is 
concerned  only  with  producing  two  kinds  of  ions  from  the  elements. 
True,  these  ions,  if  they  united,  would  give  the  product  shown  in  the 
equation  (ZnBr2),  but  the  union,  if  it  ever  occurred,  would  be  without 
electrical  effect.  It  is  clear  that,  since  there  is  sodium  chloride  (or 
some  other  ionogen)  in  all  parts  of  the  cell,  molecules  are  ionizing,  and 
ions  are  combining,  continually,  throughout  the  whole  system. 
Thus,  on  the  left  some  zinc  chloride  molecules  are  formed  and  on 
the  right  some  sodium  bromide  molecules,  and  eventually,  near  the 
center,  some  zinc  bromide  molecules.  But  these  reactions  occur 
in  every  solution  containing  ionogens,  without  giving  any  current. 
In  a  cell,  the  only  reactions  which  contribute  to  the  current  are 
those  taking  place  at  the  surfaces  of  the  poles. 

A  Displacement  Cell.  —  In  a  similar  way,  a  cell  using  metallic 
zinc  and  cupric  sulphate  solution  may  be  arranged  (Fig.  151).  The 
zinc  forms  one  pole,  and  the  cupric  sul- 
phate solution  must  be  placed  on  the  other 
side  of  the  partition.  For  inactive  mate- 
rials, a  plate  of  copper  or  of  some  metal  be- 
low copper  in  the  activity  series  may  be 
used,  and  any  solution  (such  as  zinc  chlo- 
ride solution)  which  will  interact  neither 
with  the  zinc  nor  with  the  cupric  sulphate. 

In  following  the  operation  of  the  cell, 
we  may  start  at  either  pole.  Thus,  the 
zinc  gives  zinc-ion  Zn°  — >  Zn"1"1"  +20. 
The  wire  becomes  negatively  charged. 
The  cupric-ion  is  discharged  on  the  other 
pole  Cu++— »  Cu°  -f  2®,  rendering  it 
positive.  All  the  positive  ions  in  the 
cell  migrate  towards  the  right  pole  (cathode).  All  the  negative  ions 
migrate  towards  the  left  pole  (anode),  since  positive  ions  are  being 
formed  on  the  left  and  are  disappearing  on  the  right. 

When  bromine  displaces  iodine,  the  cell  may  be  arranged  as 
in  Fig.  152.  The  iodine  liberated  dissolves  in  the  potassium  iodide 
solution  and,  with  starch  emulsion  present,  its  formation  can  be 
detected  in  a  few  seconds. 


4— SO«+Cu 


Pos.  iQns  -*•         Neg.  ions-*- 
Fia.  151. 


792 


INORGANIC   CHEMISTRY 


An  Oxidation  Cell.  —  The  arrangement  whereby  stannous-ion 
Sn++  is  oxidized  by  chlorine-water  to  stannic-ion  Sn++++  is  shown  in 
Fig.  153.  The  chlorine  Cl°  encountering  the  pole  becomes  nega- 
tively charged,  leaving  the  pole  positive.  This  positive  charge  is 


/fey.  ions-*         Neg.  ions-*- 
FIG.  152. 


Pos.  ions-*         Neg.  ions- 
Fia.  153. 


shared  by  the  whole  conducting  wire  and,  at  the  other  pole,  furnishes 
the  positive  electricity  required  to  raise  the  charge  of  each  tin  ion 
from  Sn++  to  Sn++++.  Only  the  tin  ions  which  touch  the  pole  can 
acquire  the  charge. 

Facts  Concerning  all  Cells.  —  If  the  wire  is  disconnected,  the 
progress  of  the  chemical  action  is  stopped,  although  the  difference  in 
potential  remains.  The  charge  conferred  upon  a  pole,  such  as  that 
from  the  cupric  ion  (Fig.  151),  must  be  conducted  away,  before 
additional  charges  will  be  transferred  to  it. 

If  a  glass  partition  is  substituted  for  a  porous  one,  the  cell  ceases 
to  generate  electricity.  The  partition  must  permit  the  transmi- 
gration of  the  ions,  which  is  a  necessary  part  of  the  operation  of 
the  cell. 

When  the  circuit  is  closed,  the  changes  described  go  on  until 
one  of  the  active  materials  is  exhausted  —  for  example,  until  all 
the  cupric-ion  has  been  deposited  as  copper,  or  until  all  the  zinc 
has  been  consumed. 

The  quantity  of  electricity  produced  is  96,504  coulombs  for  each 
equivalent  weight  of  the  active  materials  transformed,  e.g.,  for 
every  65.4/2  g.  of  zinc  consumed.  The  rate  at  which  the  elec- 


ELECTROMOTIVE  CHEMISTRY  793 

tricity  is  produced  is,  in  general,  greater  the  larger  the  area  of  the 
poles.  The  amperage  of  a  single  cell  is,  in  general,  very  low. 

The  E.M.F.  of  the  cell  is  not  changed  by  altering  the  size  or 
shape  of  the  poles,  or  by  using  more  or  less  of  the  solutions.  It 
is  affected  by  any  change  in  the  qualities  of  the  active  materials, 
however.  Changing  the  concentration  (see  p.  794)  of  the  cupric- 
ion  (Fig.  151)  or  of  the  bromine-water  (Fig.  152)  has  an  immediate 
effect.  So  has  substituting  one  active  metal  for  another  (see  p. 
796),  as  magnesium  for  zinc  (Fig.  150).  Even  hammering  the 
metal,  thus  making  it  denser,  has  a  slight  effect. 

Single  Potential  Differences  Produced  by  the  Metals.  —  If 

we  reconsider  the  cells  described,  we  shall  see  that  there  are  really 
two  chemical  actions  in  each  cell  and  that  these  are  to  some  extent 
independent.  We  can  leave  the  zinc  (Fig.  150)  constant,  and  change 
the  concentration  of  the  bromine  or  even  substitute  chlorine  or  iodine 
for  the  latter.  On  the  other  hand,  we  can  leave  the  bromine-water 
constant,  and  exchange  the  zinc  for  some  other  active  metal.  Thus, 
the  E.M.F.  of  every  cell  is  really  the  resultant  of  two  effects.  Now, 
these  effects  can  be  measured,  separately. 

If  we  place  zinc  in  a  solution  of  zinc  chloride,  we  find  that  there 
is  at  once  a  difference  in  potential  between  the  metal  and  the  solu- 
tion! The  metal  has  an  individual  tendency  to  become  ionic  — 
a  sort  of  solution  pressure  *  —  and  to  form  a  few  ions,  thus  making 
the  liquid  positive  and  the  metal  negative.  In  reality,  it  is  the 
tendency  of  the  atoms  of  the  metal  to  give  up  electrons  (p.  354), 
e.g.,  Zn  —  2e  =  Zn++,  which  is  being  observed.  On  the  other 
hand,  the  ions  have  a  tendency  to  deposit  themselves,  and  a  few 
may  be  deposited  (taking  up  their  electrons  and  becoming  neutral), 
rendering  the  pole  positive  and  the  solution  'negative.  If  the  former 
tendency  (the  tendency  to  give  up  electrons)  is  the  stronger  of 
the  two  (the  more  active  metals),  then  a  difference  in  potential  is 
produced,  with  the  solution  positive.  If  the  latter  tendency  is  the 
stronger  (less  active  metals)  the  solution  is  observed  to  be  negative. 
Since  raising  the  concentration  of  the  metal-ion  will  increase  the 

*  The  charge  of  electricity  is  apt  to  interfere  with  the  ready  acceptance  of 
this  idea.  If  it  is  remembered  that  the  ionic  form  of  an  element  is  simply  an 
allotropic  modification,  with  a  different  amount  of  available  energy,  the  difficulty 
disappears.  In  the  ionic  allotrope  the  free  energy  is  sometimes  greater  (nickel 
to  gold)  and  sometimes  less  (potassium  to  cobalt)  than  in  the  free  element  (see 
below). 


794  INORGANIC  CHEMISTRY 

tendency  to  deposition,  and  vice  versa,  it  is  customary  to  take  as  the 
standard  solution,  for  this  purpose,  one  in  which  the  concentration 
of  the  metal-ion  is  normal  (N).  The  ions  of  the  foreign  salt,  if 
such  a  salt  was  introduced  at  first,  need  not  be  considered.  In  the 
following  table,  the  sign  preceding  the  number  is  that  of  the  charge 
of  the  solution. 

POTENTIAL  OF  N  SOLUTIONS  IN  CONTACT  WITH  METALS 
(E.-M.  SERIES) 

K  (+2.6)  Fe  (Fe++)  +0.2  Bi  -0.63? 

Na  (+2.4)  Cd  +0.16  Sb  -0.71? 

Ba  (+2.6)  Co  +0.05  Hg  (Hg+)  -0.99 

Sr  (+2.5)  Ni  -0.02?  Pd  -1.03? 

Ca  (+2.4)  Pb  -0.12  Ag  -1.04 

Mg  +1.3  Sn  (Sn++) -0.14  Pt  -1.10? 

Al  +1.0  H  -0.24  Au  -1.7? 

Mn  +0.8  As  -0.53 

Zn  +0.5  Cu(Cu++)-0.58 

Thus,  opposite  Mg  we  find  +1.3.  This  means  that  when  a 
piece  of  magnesium  is  placed  in  a  solution  of  a  salt  of  magnesium, 
containing  normal  concentration  of  Mg^,  the  solution  is  positively 
charged  (the  metal  negatively)  and  the  difference  in  potential  is  1.3 
volts.  With  silver  in  a  solution  of  a  salt  of  silver,  containing  normal 
concentration  of  silver-ion,  the  solution  is  negative  and  the  difference 
in  potential  is  —1.04  volts. 

For  a  hydrogen  pole,  a  piece  of  palladium  saturated  with  hydro- 
gen at  760  mm.  (p.  124)  is  used.  The  values  for  the  metals  which 
decompose  water  with  ease  cannot  be  observed,  and  so  calculated 
values  are  given  in  parentheses.  An  interrogation  point  indicates 
that  the  value  is  uncertain. 

Calculation  of  the  Potential  with  Varying  Concentrations. 

—  These  facts  enable  us  to  state  in  more  definite  terms  the  formu- 
lative  hypothesis  foreshadowed  above.  It  was  first  put  forward  by 
Nernst. 

Every  metal  has  a  certain  solution  tension,  or  pressure,  tending  to 
drive  it  into  solution  (in  ionic  form,  of  course,  since  it  is  not  soluble 
otherwise).  The  value  of  this  pressure  becomes  rapidly  less  as  we 
pass  through  the  series  from  magnesium  to  gold.  If  the  ions  of  the 
same  metal  are  already  present,  they  tend  to  give  up  their  electrical 


ELECTROMOTIVE  CHEMISTRY  795 

charges  and  deposit  themselves  upon  the  metal.  These  two  tend- 
encies oppose  one  another,  just  as  solution  pressure  and  diffusion 
pressure  oppose  one  another  in  the  ordinary  process  of  dissolving 
any  substance  (p.  186).  When  the  tendency  of  the  ions  to  deposit 
themselves  is  the  greater  of  the  two,  a  very  minute  excess  of  dep- 
osition over  solution  occurs,  and  thus  the  solution  has,  as  a  whole, 
a  negative  charge  (having  lost  some  positive  ions),  and  the  metal 
has  a  positive  charge  (having  acquired  it  from  the  deposit  of  a  few 
ions).  This  is  the  case  with  gold  and  the  metals  as  far  up  the  list 
as  nickel.  When,  on  the  other  hand,  the  solution  pressure  of  the 
metal  is  the  greater  of  the  two,  the  solution  acquires  a  very  slight 
excess  of  positive  ions,  and  is,  therefore,  positively  charged  when 
compared  with  the  metal.  This  is  the  case  from  potassium  down 
krcobalt. 

The  measure  of  the  "tendency  of  the  ions  to  deposit  themselves" 
is  simply  the  diffusion  pressure  of  the  metal-ions.  We  perceive  this 
to  be  the  case,  for,  when  we  take  a  stronger  solution  of  the  salt,  and 
therefore  an  increased  diffusion  pressure  of  the  ions,  an  instant  effect 
is  produced.  The  solution  becomes  less  positive,  or  more  negative, 
as  the  case  may  be.  Evidently  the  solution  pressure  of  each  of  the 
metals  near  to  cobalt  and  nickel  is  almost  exactly  balanced  by  the 
diffusion  pressure  of  a  normal  solution  of  the  ions  composed  of  the 
same  metal.  This  pressure,  for  a  univalent  metal,  is  22.4  atmospheres 
(p.  333),  and  for  a  bivalent  metal  11.2  atmospheres.  The  metals 
above  nickel  have  solution  pressures  higher  and  higher  above  this 
norm;  those  below  cobalt  have  solution  pressures  farther  and  farther 
below  it.  The  effect  of  changing  the  diffusion  pressure  is  independent 
of  the  particular  substances  used,  and  depends  only  on  the  valence. 
When  the  concentration  of  the  metal-ion  becomes  ION,  0.058  volts 
must  be  subtracted  (algebraically)  from  the  potential  (see  above 
table)  of  the  liquid,  if  the  metal-ion  is  univalent.  If  it  is  n-valent, 
0.058/n  must  be  subtracted.  When  the  solution  is  0.1  AT,  0.058/n 
volt  must  be  added;  when  it  is  O.OL/V,  2  X  0.058/n  must  be  added, 
and  so  forth.  Thus,  zinc  with  decinormal  zinc-ion  gives  +0.5  + 
0.058/2  =  0.53  volts,  approximately;  silver  with  centinormal  silver- 
ion  gives  — 1.04  +  (2  X  0.058)  =  —0.924  volts,  approximately.  And, 
in  general,  if  c  be  the  equivalent  concentration  of  the  metal-ion  in 
the  liquid  under  consideration,  and  TTC  the  electrical  potential  of  that 
liquid,  while  TTN  is  the  potential  of  the  liquid  containing  N  metal-ion, 

,  0.058  ,     1 

1TC   =  7J>H log-' 

n  c 


796 


INORGANIC  CHEMISTRY 


Applications:  E.M.F.  of  a  Displacement  Cell.  —  When, 
now,  a  cell  with  two  poles  and  two  metal-ions  is  set  up,  we  can  tell 
from  the  above  table  what  the  difference  in  potential  between  the  two 
poles  will  be.  We  may  regard  the  two  systems  —  the  anodic  and 
cathodic  —  as  working  against  each  other.  Each  metal  tends  to 
project  its  ions  into  the  solution  and  to  generate  a  positive  current 
in  the  liquid  and  a  negative  one  in  the  wire.  If  both  solutions  are 
normal,  or,  in  general,  of  equal  equivalent  concentration,  the  relative 
solution  pressures  of  the  metals  decide  the  direction  of  the  resultant 
current,  and  its  magnitude  will  be  the  difference  of  the  two  effects. 
Thus,  the  values  for  the  following  pairs  will  be : 

Zn-Cd++,  +0.5    -  (+0.16)  =  +0.34,  Zinc  the  negative  pole. 
Cd-Cu++,  +0.16  -  (-0.58)  =  +0.74,  Cadmium  the  negative  pole. 
Zn-Cu++,  +0.5    -  (-0.58)  =  +1.08,  Zinc  the  negative  pole. 

The  Daniell  or  gravity  cell  (Fig.  154)  represents  the  last  of  these 
three  combinations.  The  copper  pole  is  at  the  bottom,  and  the  zinc 
plate  is  suspended  above  it.  The  cell  is  charged 
with  a  dilute  solution  of  sodium  chloride,  and 
blue  vitriol  crystals  are  thrown  in  and  dissolve. 
So  long  as  the  contents  are  not  disturbed,  the 
solutions  require  no  porous  septum  to  keep  them 
apart.  It  is  true  that,  when  the  current  is  not 
being  used,  and  the  cell  is  not  working,  the  cupric 
sulphate  diffuses  upwards.  During  the  time  that 
the  circuit  is  closed,  however,  the  effects  of  diffu- 
sion are  nullified  by  the  migration  of  the  cupric- 
ion  away  from  the  zinc  and  towards  the  positive 
pole.  The  actual  E.M.F.  of  this  cell  is  not  ex- 
actly that  calculated  above  for  normal  solutions,  because  the  cupric 
sulphate  is  in  saturated  solution,  and  the  concentration  of  the  zinc- 
ion  varies,  starting  at  zero  and  increasing  as  the  cell  is  used.  It  is, 
however,  a  little  over  1  volt. 

The  cell  Zn-H+  (+0.5  -  (-0.24)  =  0.74  volts)  works  with- 
out a  septum,  provided  the  direct  action  of  the  zinc  on  the  acid  is 
minimized  by  adequate  amalgamation  with  mercury.  It  gives  a 
very  inconstant  electromotive  force,  however,  because  the  platinum 
plate  used  as  the  cathode  becomes  covered  with  bubbles  of  hydrogen, 
and  so  the  internal  resistance  of  the  cell  is  greatly  increased.  The 
polarization  (p.  360)  also  diminishes  the  electromotive  force.  These 


FIG.  154. 


ELECTROMOTIVE  CHEMISTRY  797 

difficulties  are  remedied,  and,  in  fact,  a  great  increase  in  the  E.M.F. 
of  the  cell  is  effected,  by  surrounding  the  cathode  with  an  oxidizing 
agent  which  shall  convert  the  hydrogen  into  water.  The  energy 
obtainable  is  thus  that  of  a  strong  oxidizing  agent  on  zinc,  and  not 
merely  that  of  an  acid.  In  the  Bunsen  cell  the  cathode  is  a  carbon 
block  surrounded  by  concentrated  nitric  acid.  In  the  dichromate 
battery  it  is  a  carbon  block  with  chromic  acid.  Each  of  these  cells 
gives  an  E.M.F.  of  1.9  volts.  In  the  Leclanche  cell  the  cathode  is  a 
mixture  of  carbon  and  manganese  dioxide,  and  the  fluid  is  a  solution 
of  ammonium  chloride  from  which  the  zinc  displaces  hydrogen.  The 
dioxide,  being  solid,  oxidizes  the  hydrogen  slowly,  and  the  cell  can 
be  used  for  only  a  few  minutes  at  a  time  without  becoming  polarized. 
The  E.M.F.  is  1.48  volts.  Dry  cells  are  of  the  same  nature,  but  con- 
tain a  porous  solid  which  holds  the  liquid  by  capillary  forces  (for 
Accumulators,  see  under  Lead). 

A  cell  is  thus  an  engine  for  the  direct  transformation  of  chemical 
into  electrical  energy,  just  as  a  steam-engine  transforms  chemical 
energy,  by  several  stages,  it  is  true,  into  mechanical  energy.  The 
cell  is  driven  by  pressure-differences  in  the  materials  in  and  around 
the  two  poles. 

The  Weston  Standard  Cell  contains  a  pole  of  mercury  in  a  satu- 
rated solution  of  mercurous  sulphate,  and  cadmium  in  contact  with 
saturated  cadmium  sulphate  solution.  For  normal  solutions,  the 
voltage  would  be  +0.16-  (-0.99)  =  1.15  volts.  At  20°  it  is 
1.0183  volts. 

The  Clark  Standard  Cell  contains  zinc  and  zinc  sulphate  solution 
in  place  of  the  cadmium.  With  normal  solutions  it  would  give 
+ 0.5  -  (-0.99)  =  1.49  volts.  It  actually  gives  1.434  volts. 

Single  Potential  Differences  for  Non-Metallic  Ions.  —  The 

corresponding  figures  for  the  non-metals  are: 

r    -o.78        cr   -1.59        Hsor  -2.9 

Br~   -1.32  OH-   -1.96 

O=    -1.36  S04=  -2.2 

As  before,  the  anode  potential  is  supposed  to  work  against  the  cathode 
potential  and  is  subtracted  from  it.  Hence  the  cell  with  zinc  and 
bromine-water  (p.  789)  in  presence  of  normal  concentration  of  the 
respective  ions,  gives  +0.5  —  (—1.32)  =  1.82  volts.  Similarly,  the 
cell  in  which  bromine  displaces  iodine  (p.  791)  gives  —0.78  — 
(-1.32)  =  0.54  volts. 


798  INORGANIC  CHEMISTRY 

Applications :     Electrolysis :     Discharging     Potentials.  — 

When  a  solution  of  a  salt,  such  as  cupric  chloride,  is  electrolyzed, 
copper  and  chlorine  are  liberated  at  the  two  poles.  Now,  when 
the  electrolysis  has  made  some  progress,  if  the  battery  is  taken 
out,  and  the  wires  are  joined,  a  current,  the  polarization  current, 
flows.  Evidently,  the  copper  and  chlorine  liberated  in  and  round 
the  electrodes  have  made  the  arrangement  into  a  copper-chlorine 
battery  cell.  Assuming  normal  concentrations,  the  E.M.F.  of  the 
polarization  current  is  —0.58  —  (—1.59)  =  1.01  volts.  Now  this 
counter-current  is  in  operation  during  the  whole  electrolysis.  To 
overcome  it,  and  maintain  the  electrolysis,  evidently  an  E.M.F.  of 
at  least  1.01  volts  from  the  battery  is  required.  This  is  called  the 
discharging  potential  for  cupric  chloride. 

Again,  a  tin-chlorine  cell  produces  —0.14  —  (—1.59)  =  1.45  volts, 
and  this  E.M.F.  will  just  suffice  to  electrolyze  tin  chloride.  Similarly, 
hydrochloric  acid  will  require  at  least  —0.24  —  (—1.59)  =  1.35 
volts,  and  zinc  sulphate  0.5  -  (-2.2)  =  2.7  volts. 

Oxygen  acids  like  sulphuric  acid  show  a  trace  of  decomposition  at 
1.12  volts  (=  -0.24  -  (-1.36)),  and  a  noticeable  but  still  small 
decomposition  at  1.72  volts  (=  -0.24  -  (-1.96)),  due  to  the  H+ 
and  O=  and  the  H+  and  OH~  respectively.  But  it  is  only  when  the 
E.M.F.  reaches  the  values  for  H+  and  S04=,  and  H+  and  HS04~, 
namely,  1.96  and  2.66  volts,  that  rapid  electrolysis  begins.  This 
observation  answers,  incidentally,  the  question  whether  in  the  so- 
called  "electrolysis  of  water,"  when  dilute  sulphuric  acid  is  used,  it 
is  the  water  or  the  acid  that  is  decomposed.  The  H+  and  OH~~  de- 
composition at  1.72  volts  is  very  slight,  because  of  the  small  concen- 
tration of  the  OH~,  and  a  lens  is  required  for  its  recognition.  The 
more  vigorous  action  resulting  from  the  discharge  of  SO4=  and  HS04~ 
by  the  use  of  2-3  volts  is  therefore  the  one  invariably  used. 

In  view  of  the  foregoing  facts,  it  is  probably  most  correct  to  say 
that  when  dilute  sulphuric  acid  is  electrolyzed,  e.g.,  as  a  lecture 
experiment,  the  oxygen  liberated  at  the  anode  comes  mainly  from  a 
secondary  interaction  of  the  discharged  material  of  the  anions  with 
the  water  (p.  121).  A  minute  proportion  of  the  oxygen  in  such  an 
experiment  does  arise  from  primary  electrolysis  of  the  water,  but  this 
effect  of  the  current  is  in  itself  too  slight  to  be  visible  at  a  distance. 
When,  on  the  other  hand,  the  solution  electrolyzed  contains  a  salt  of 
sodium,  and  hydrogen  is  liberated  at  the  cathode,  this  gas  must  be 
regarded  as  coming  chiefly  from  primary  electrolysis  of  the  water. 
The  discharging  potential  for  sodium  chloride  should  be  +2.4  — 


ELECTROMOTIVE  CHEMISTRY  799 

(—1.59)  =  3.99  volts,  and  with  the  help  of  a  mercury  cathode  a 
sodium  amalgam  is  easily  obtained  (p.  667).  But  it  will  be  found 
that,  with  platinum  electrodes,  hydrogen  and  chlorine  are  liberated 
freely  from  a  solution  of  salt  by  a  current  of  little  more  than  half  the 
above  mentioned  E.M.F.  The  positive  electricity  is  carried  in  the 
liquid  mainly  by  the  very  numerous  sodium  ions.  But,  when  these 
ions  reach  the  cathode,  the  potential  difference,  being  insufficient 
to  discharge  the  sodium-ion,  liberates  the  hydrogen-ion  of  the  water 
instead.  Thus  the  accumulating  hydroxide-ion  of  the  water,  and  the 
sodium-ion  arriving  by  migration,  together  constitute  the  sodium 
hydroxide  which  is  another  product  of  this  electrolysis.  With  high 
E.M.F.  and  sufficient  current  density,  sodium-ion  is  doubtless  actu- 
ally discharged,  and  in  that  case  a  part  of  the  hydrogen  liberated  is 
furnished  by  the  interaction  of  the  metal  with  the  water. 

The  ordinary  chemical  behavior  of  the  halogens  accords  with  the 
order  of  their  potential  differences.  Bromine  displaces  iodine,  and 
chlorine  displaces  both  (p.  285) .  Chlorine,  however,  does  not  displace 
easily  perceptible  amounts  of  oxygen  from  water,  because  of  the  small 
concentration  of  the  0~  (obtained  by  secondary  ionization  of  the 
OH~).  The  oxygen  freely  liberated  in  sunlight  comes  from  the  de- 
composition of  the  HC10  (p.  223).  Fluorine,  however,  which  would 
show  a  potential  difference  greater  than  that  of  the  much  more  plen- 
tiful OH~,  displaces  oxygen  vigorously  by  discharging  this  ion. 

Applications:  Electrolytic  Refining.  —  The  electrolytic  proc- 
ess of  refining  copper  (read  p.  747)  can  now  be  more  easily  under- 
stood. Both  electrodes  are  made  of  copper,  and  the  solution  contains 
cupric  sulphate.  There  is,  therefore,  no  difference  in  potential  be- 
tween the  plates,  except  a  very  small  one,  due  to  the  fact  that  one 
plate  is  pure  copper  and  the  other  impure.  Hence  a  very  slight 
E.M.F.,  sufficient  to  overcome  the  difference  just  mentioned,  and  to 
overcome  the  friction  of  the  moving  ions,  is  all  that  is  required,  and 
0.5  volts  is  sufficient. 

As  regards  the  resulting  purification,  the  anode  of  crude  copper, 
which  is  being  consumed,  contains,  besides  copper,  small  amounts 
of  less  active  metals  like  silver  and  gold,  and  of  more  active  metals 
like  zinc.  So  far  as  the  more  active  metals  are  concerned,  the 
cell  is  like  one  with  zinc  and  cupric  sulphate  (p.  791).  It  would 
run  by  itself,  without  any  outside  current,  and  would  actually  gen- 
erate a  current.  Hence  the  active  metals  become  ionic  easily,  and 
displace  cupric-ion  from  the  solution.  The  less  active  metals,  on 


800  INORGANIC  CHEMISTRY 

the  other  hand,  are  not  required  for  the  transference  of  the  electricity, 
since  a  great  excess  of  the  more  active  copper  is  available.  They 
also  require  a  larger  E.M.F.  for  their  ionization  than  does  copper. 
Hence,  they  remain  as  metals,  and  drop  to  the  bottom  of  the  cell 
(sludge)  as  the  anode  of  crude  copper  wears  away. 

Applications:  Couples.  —  The  fact  that  metallic  zinc  will  dis- 
place hydrogen-ion  from  an  acid,  or  cupric-ion  from  cupric  sulphate 
solution  can  now  be  explained.  The  more  active  metals  are  the  ones 
which  have  the  greatest  tendencies  to  become  ionic.  Each  will  de- 
prive the  ions  of  a  metal  below  it  in  the  list  of  their  electric  charges : 

Zn°  +  2H+  ->  Zn++  +  H2°T. 
Zn°  +  Cu++  rr>  Zn++  +  Cu°| . 

Now  we  have  noted  the  facts  (pp.  119,  769)  that  contact  with  a 
platinum  wire,  or  the  presence  of  impurities  (other  metals)  in  the 
zinc,  will  hasten  its  action.  Pieces  of  two  metals  in  contact  with 
one  another  constitute  a  couple.  With  zinc  and  platinum  in  an 
acid,  a  current  is  set  up,  like  that  of  a  short  circuited  cell.  The 
zinc  becomes  negative,  the  platinum  positive,  and  the  hydrogen  is 
liberated  upon  the  platinum.  This  facilitates  the  action  because, 
when  the  platinum  is  absent,  and  the  hydrogen  gas,  in  bubbles,  is 
liberated  on  the  surface  of  the  zinc,  this  surface  is  only  partly  in  con- 
tact with  the  acid  (H+),  and  so  the  liberation  of  the  hydrogen  is 
slower. 

Galvanized  iron  is  also  a  couple.  When  rain  (dilute  carbonic 
acid)  falls  upon  it,  the  zinc,  being  the  more  active  metal  (p.  794), 
is  the  anode  and  tends  to  become  ionized  (forming  the  carbonate). 
The  iron  is  the  cathode  and  is  not  affected.  The  carbonate,  how- 
ever, forms  a  closely  adhering  coating  on  the  zinc,  and  so  but  little 
of  this  metal  is  actually  consumed,  and  the  material  is  therefore 
durable.  On  the  other  hand,  a  sheet  of  iron,  without  the  zinc  coat- 
ing, gives  ferrous  carbonate  which  is  easily  oxidized  to  ferric  hydroxide 
(a  base  too  weak  to  give  a  carbonate).  This  forms  a  brittle,  porous 
layer  which  does  not  mechanically  protect  the  surface  from  further 
action,  and  so  the  iron  is  finally  all  oxidized.  Tin-plate  (tin  on 
iron,  a  couple)  is  not  attacked  so  long  as  the  layer  of  tin  is  nowhere 
broken.  But  damaged  tin-plate  rusts  rapidly.  There,  the  iron  is 
the  more  active  metal  (p.  794)  and  forms  carbonate  and  then  hydrox- 
ide continuously,  while  the  tin  remains  unaffected. 


ELECTROMOTIVE  CHEMISTRY 


801 


Rusting.  —  It  seems  to  be  generally  overlooked  that  two  factors  are  con- 
cerned in  determining  extent  to  which  the  rusting  or  tarnishing  of  a  metal 
will  proceed.  Other  things  being  equal,  it  should  depend  upon  the  order  of 
activity  of  the  metals,  since  it  is  a  question  of  action  upon  surface  moisture  con- 
taining carbonic  acid,  with  atmospheric  oxygen  as  an  assistant  when  the  metal 
does  not  liberate  elementary  hydrogen  (e.g.,  Copper,  p.  745) ..  But  this  applies  only 
to  the  initiation  of  the  action.  The  physical  qualities  of  the  product  then  deter- 
mine whether  the  rusting  shall  be  superficial  or  shall  go  deeper.  Magnesium 
(p.  765)  and  iron  (p.  800)  give  carbonates,  and  the  latter  eventually  a  brittle, 
scaly  hydrated  oxide.  In  both  cases,  the  porous  product  harbors  moisture  and 
thus  promotes  further  rusting.  Aluminium  is  more  active  than  is  iron,  but  the 
horny,  gelatinous,  closely  adherent  hydroxide  first  formed  protects  the  surface 
from  further  action.  The  same  is  true  of  zinc  and  tin. 

Concentration  Cells.  —  If  two  rods  of  a  metal  (e.g.,  tin)  are 
placed  together  in  the  same  solution  of  a  salt  of  the  metal  (e.g., 
stannous  chloride  SnCl2),  there  is  no  difference  in  potential,  because 
the  state  of  both  poles  is  in  all  respects  the  same,  and  no  current 
flows  when  they  are  connected  by  a  wire.  The  two  poles  and  their 
solutions  are  here  alike  and  —0.14  —  (—0.14)  =  0.  But  if  the 


«-8n-H- 


<-Sn-H-+2CJ- 


Pos.  ions-*- 


Afop.  ions+r- 


FIG.  155. 


FIG.  156. 


solution  round  one  pole  is  diluted  to  N/1Q  concentration,  the  poten- 
tial at  that  pole  becomes  at  once  — 0.14  +  0.058/2  =  —0.11  volts, 
approximately,  and  a  current  is  set  up  (Fig.  155).  The  tendencies 
of  the  metallic  tin  to  form  ions  are  equal,  but  the  pressures  of  the 
stannous  ions  are  different,  and  so,  when  the  circuit  is  closed,  stan- 
nous ions  are  discharged  on  the  tin  pole  in  the  more  concentrated 


802  INORGANIC   CHEMISTRY 

solution,  forming  long  crystals  of  tin,  and  tin  in  equal  amount  from 
the  pole  in  the  dilute  solution  becomes  ionic. 

Figure  156  shows  the  simplest  arrangement,  where  the  more  con- 
centrated, denser  liquid  is  below,  and  one  rod  of  tin,  passing  through 
both  layers,  furnishes  at  once  the  two  poles  and  the  connection. 
The  chloride-ion  migrates  through  the  solution,  in  a  direction  opposite 
to  that  taken  by  the  tin  ions,  and  thus  passes  upwards  into  the 
dilute  solution  to  balance  the  fresh  tin  ions  that  are  continuously 
formed.  All  change  ceases  when  the  concentrations  have  become 
equalized. 

The  concentration  cell  is  instructive  because  it  shows  that  the 
order  of  the  metals  in  the  electromotive  series  is  not  determined  by 
the  metal  alone,  but  also  by  the  concentration  of  the  solution.  The 
order  of  the  metals  in  the  electromotive  series  is  therefore  subject  to 
variation.  An  extreme  case  of  this  occurred  in  a  recent  chapter. 
Zinc  displaces  copper  from  a  solution  of  a  cupric  (or  cuprous)  salt, 
and  any  but  a  prodigiously  dilute  solution  will  show  the  effect. 
But  a  solution  containing  cuprocyanide-ion  Cu(CN)2~  has  precisely 
this  very  minute  concentration  of  copper  ions  which  will  turn  the 
scale.  Hence,  zinc  will  not  displace  copper  from  this  solution 
(p.  744).  On  the  contrary,  copper  will  displace  zinc  from  a  solution 
of  a  salt  of  the  latter  containing  excess  of  potassium  cyanide,  and 
therefore  the  complex  salt  K.Zn(CN)3. 

The  law  which  formulates  the  relation  between  the  two  concen- 
trations and  the  E.M.F.  produced  being  known  (p.  795),  it  is  possible 
to  use  the  concentration  cell  for  measuring  solubilities  of  insoluble 
salts.  Thus,  we  cannot  easily  measure  the  solubility  of  silver  chlo- 
ride by  the  ordinary  method  (p.  180),  because  evaporation  of  the 
solution  may  leave  a  larger  mass  of  impurities,  derived  from  solution 
of  the  glass,  than  of  dissolved  silver  chloride.  Hence,  we  use  two 
poles  of  silver,  place  one  in  normal  silver  nitrate  solution  and  the 
other  in  saturated  silver  chloride  solution  (with  excess  of  the  solid), 
measure  the  difference  in  potential,  and  calculate  the  ratio  of  the 
concentrations  of  silver-ion  in  the  two  solutions.  The  absolute 
value  of  that  in  the  silver  nitrate  solution  is  known,  and  so  the 
absolute  value  of  the  Ag+  concentration  in  the  silver  chloride 
solution  can  be  found.  Since  silver  chloride  is  a  salt  (p.  369),  it 
is  very  highly  ionized  in  s6  dilute  a  solution,  and  the  molecular 
concentration  of  silver-ion  is  practically  equal  to  the  total  molec- 
ular concentration  of  silver,  and  therefore  of  silver  chloride  in  the 
liquid. 


ELECTROMOTIVE  CHEMISTRY  803 

The  Factors  of  Energy.  —  We  have  seen  that  the  amount  of  a 
given  supply  of  electrical  energy  is  described  by  two  factors,  the 
E.M.F.  and  the  quantity  of  electricity,  and  that  the  weight  of  mate- 
rial, which,  by  its  influence,  undergoes  a  given  chemical  change,  is 
proportional  solely  to  this  second  factor.  On  the  other  hand,  the 
question  whether  the  supply  of  energy  can  initiate  the  change  at  all  de- 
pends on  the  magnitude  of  the  first  factor  alone  (p.  798) .  The  total 
amount  of  available  energy  does  not  influence  the  result  if  the  E.M.F. 
is  not  above  a  certain  minimum,  which  differs  from  case  to  case.  Now 
the  same  is  true  of  other  kinds  of  energy.  The  quantity  of  each  may 
be  expressed  as  the  product  of  an  intensity  factor  and  a  capacity 
factor.  The  magnitude  of  the  former  determines  whether  the  energy 
can  be  transferred  or  transformed  or  not.  Heat  energy,  no  matter 
how  much  of  it  is  at  hand,  can  neither  flow  nor  be  transformed  into 
work  unless  the  source  is  at  a  higher  temperature  (intensity  factor) 
than  the  surroundings.  A  head  of  water  will  do  work  only  when  it 
is  connected  with  a  receptacle  at  a  lower  level.  It  is  the  pressure 
of  the  water  that  determines  its  availability.  The  E.M.F.  is  the 
intensity  factor  of  electrical  energy. 

Now  we  may  presume  that  chemical  energy  can  also  be  expressed 
by  two  factors.  One  of  these,  the  capacity  factor,  must  be  pro- 
portional to  the  quantity  of  material,  in  other  words,  to  the  number 
of  chemical  equivalents.  The  other  is  the  chemical  potential  (inten- 
sity factor).  A  chemical  change  which  does  not  take  place  on  a 
small  scale  will  not  take  place  when  more  material  is  used,  provided 
the  relative  amounts  of  the  interacting  substances  and  the  conditions 
remain  unchanged.*  We  have,  in  fact,  been  assuming  all  along  that 
this,  the  capacity  factor,  is  not  the  most  significant  one.  But  we 
have  devoted  ourselves  to  noting  such  things  as  these:  that  chlorine 
will  displace  bromine,  and  therefore  has  the  higher  potential  of  chemi- 
cal energy;  that  magnesium  reduces  sand,  while  hydrogen  does  not, 
and  that  magnesium  is  therefore  a  more  active  reducing  agent; 
and  that  hypochlorous  acid  will  oxidize  indigo,  while  free  oxygen  will 
not,  and  is  therefore  a  more  powerful  oxidizing  agent.  When  we 
were  comparing  degrees  of  activity,  therefore,  we  were  really  trying 
to  describe  the  relative  potential  of  the  chemical  energy  in  all  sorts  of 
substances.  At  present,  the  state  of  the  science  permits  this  to  be 
done  in  most  cases  in  a  rough  fashion  only. 

*  Change  in  concentration,  a  condition,  however,  does  affect  activity,  and 
therefore  modifies  the  chemical  potential. 


804  INORGANIC  CHEMISTRY 

Applications:  Measurement  of  Affinity.  —  Since  the  capac- 
ity factor  of  chemical  energy  is  proportional  to  the  number  of 
equivalent  weights  transformed,  and  the  capacity  factor  of  electrical 
energy  is  proportional  to  the  same  thing  (Faraday's  law),  it  follows 
that  the  intensity  factor  of  the  chemical  energy  (the  chemical  poten- 
tial) in  a  given  substance  undergoing  a  given  change,  must  be  pro- 
portional to  the  corresponding  factor  (the  E.M.F.)  of  the  electrical 
energy  produced  when  the  same  change  takes  place  in  a  suitable  cell. 
Thus  the  activities  of  the  metals,  expressed  in  volts  (p.  794),  are 
accurate  figures  for  the  relative  affinities  of  the  metals,  so  far  at 
least  as  ionic  actions  are  concerned.  In  point  of  fact,  they  express 
also  the  approximate  affinities  of  the  metals  in  other  actions  (p.  404) 
as  well.  Again,  by  using  different  oxidizing  agents  in  place  of  the 
chlorine-water  (p.  792)  and  noting  the  differences  in  potential,  we  can 
obtain  numbers  representing  the  relative  activities  of  various  oxidiz- 
ing agents  towards  oxidizable  ions.  The  potential  differences  de- 
scribed above  are  therefore  often  much  more  significant  than  are 
the  results  of  thermochemical  measurements,  for  the  latter  attempt 
to  give  only  the  gross  quantity  of  chemical  energy  (in  terms  of  the 
equivalent  amount  of  heat  energy),  and  not  the  values  of  the  factors. 
The  potential  differences  come  nearer,  therefore,  to  giving  us  absolute 
values  for  chemical  activity  than  do  any  other  data  we  possess. 

As  we  have  noted  before  (p.  100),  in  spite  of  the  enormous  range 
of  temperature  at  our  disposal,  extending  to  a  point  far  above  2500° 
in  the  electric  furnace,  there  are  many  substances  for  whose  decom- 
position a  sufficient  potential  of  heat  energy  is  not  available.  On  the 
other  hand,  amongst  substances  that  are  capable  of  furnishing  an 
electrolyte,  when  dissolved  in  a  suitable  solvent  or  when  fused,  there 
are  few  that  are  not  decomposable  by  a  current  with  an  E.M.F.  of  less 
than  10  volts.  Hence  even  the  elements  which  give  the  most  stable 
compounds  and  are  the  most  difficult  to  isolate,  such  as  calcium  and 
aluminium,  are  liberated  by  electrical  methods  with  extreme  ease. 

The  idea  that  every  kind  of  energy  is  described  by  two  factors,  which  play 
different  roles,  serves  to  clarify  our  thoughts  about  many  things.  For  example,  one 
author  says:  "Valence  is  a  form  of  chemical  energy."  Now  valence  is  the  num- 
ber of  equivalents  contained  in  the  atomic  weight.  It  is  the  value  of  the  capacity 
factor,  only,  of  the  energy  in  one  atomic  weight.  It  is  not  a  form  of  energy, 
because  it  takes  no  account  of  the  intensity  factor.  Again,  the  same  author  says: 
"Chemical  energy  is  identical  with  chemical  affinity."  Yet  the  affinity  between 
a  single  atom  of  sodium  and  a  single  atom  of  chlorine  is  the  same  as  that  between 
23  kilograms  of  the  one  and  35.46  kilograms  of  the  other,  although  there  is  only  a 


ELECTROMOTIVE  CHEMISTRY  805 

trace  of  energy  produced  by  union  of  the  two  atoms  and  a  relatively  immense 
amount  by  the  union  of  the  larger  quantities.  Chemical  affinity  is  the  intensity 
factor  of  chemical  energy  only,  and  we  must  have  also  the  amount  of  material 
(capacity  factor)  before  the  whole  statement  conveys  any  information  about 
energy.  Fourteen  is  a  factor.  If  the  other  factor  is  pins,  the  amount  of  material 
is  specified,  but  very  small.  If  the  other  factor  is  elephants,  then  the  amount  of 
matter  is  much  greater.  Fourteen  by  itself  conveys  no  meaning. 

Methods  of  Measuring  Chemical  Activity.  —  The  following 
is  a  summary  of  the  methods  of  measuring  chemical  activity. 

The  thermochemical  method  (p.  98)  can  be  used  in  every  chemi- 
cal change.  But  the  heats  of  reaction  represent  the  free  energy,  and 
therefore  the  affinity,  only  when  the  heat  capacity  of  the  products  is 
equal  to  that  of  the  factors  and  no  changes  in  concentration  arise. 

For  measuring  the  activity  of  acids  in  dilute  solution,  several 
methods  have  been  mentioned:  The  speed  of  interaction  of  different 
acids  with  the  same  metal  (p.  128) ;  the  acceleration  of  the  speed  of 
hydrolysis  of  ethyl  acetate  (p.  617-8)  and  of  cane-sugar  (p.  606)  by 
different  acids;  the  amounts  of  insoluble  salts,  such  as  calcium  oxa- 
late  (p.  713),  or  zinc  sulphide,  which,  when  the  system  has  reached 
equilibrium,  are  found  to  have  been  decomposed  by  different  acids 
under  like  conditions;  the  relative  extents  of  the  hydrolysis  of  salts  of 
different  weak  acids  (p.  648) ;  the  electrical  conductivity  (p.  365)  and 
the  freezing-  and  boiling-points  of  solutions  of  acids  (pp.  336,  337). 
These  last  measure  by  physical  methods  the  same  thing  that  the 
others  determine  by  chemical  means,  namely,  the  tendency  to  ioniza- 
tion  on  which  the  activity  of  acids  depends  (p.  369.  See  also  p.  392). 

For  measuring  the  activity  of  bases  we  have:  The  relative  speeds 
of  saponification  of  esters  by  different  bases  (p»  618) ;  the  relative  ex- 
tents of  the  hydrolysis  of  salts  of  different  weak  bases  (p.  399) ;  the 
conductivity  and  the  freezing-  and  boiling-point  methods,  which 
measure  by  physical  means  the  tendency  to  ionization. 

For  measuring  the  relative  activities  of  metals  and  of  non-metals 
we  have:  The  single  potential  differences  (pp.  794,  797);  and,  for  the 
former,  the  speeds  of  interaction  of  different  metals  with  the  same 
acid  (p.  128). 

For  measuring  the  relative  activity  in  non-reversible  actions  we 
have:  The  speed  with  which  the  actions  take  place  under  like  con- 
ditions (p.  294). 

For  measuring  the  relative  activities  of  the  opposed  actions  in 
reversible  changes,  we  have:  The  concentrations  of  the  materials  re- 
maining when  equilibrium  has  been  reached  (p.  298).  The  relative 


806  INORGANIC  CHEMISTRY 

activities  in  different  reversible  changes  may  also  be  ascertained  by 
comparing  the  concentrations  in  one,  at  equilibrium,  with  those  in 
another  (cf.  p.  299). 

For  measuring  the  relative  activities  of  oxidizing  and  reducing 
agents,  we  have :  The  potential  differences  in  cells  arranged  after  the 
manner  of  the  Bunsen  and  Leclanche  cells  (p.  792). 

If  we  consider  the  whole  mass  of  phenomena,  it  must  be  admitted 
that  the  scientific  study  of  the  quantities  of  material  has  reached  a  far 
higher  level  of  exactness,  and  has  very  much  more  nearly  enveloped 
the  whole  field  covered  by  the  science,  than  has  the  study  of  relative 
activity.  Yet  it  is  evident  that  within  the  past  few  years  substantial 
advances  have  been  made  in  this  direction  also. 

Exercises.  —  1.  Make  diagrams  of  the  following  cells,  choosing 
with  care  suitable  inactive  substances  to  complete  the  arrangement: 
(a)  chlorine-water  and  aluminium;  (6)  chlorine-water  and  ferrous 
chloride;  (c)  zinc  and  dilute  sulphuric  acid;  (e)  chlorine-water  and 
potassium  bromide. 

2.  Calculate  the  E.M.F.  of  each  of  the  cells  in  Ex.  1,  assuming 
normal  solutions  to  be  present. 

3.  What  will  be  the  discharging  potentials  of  solutions  of  the 
following  substances,  assuming  N  concentrations  of  the  ions:    (a) 
manganous   chloride;     (6)    hydrogen  iodide;     (c)   ferrous  bromide; 
(e)  sodium  chloride  (hydrogen  is  liberated)? 

4.  What  weight  of  zinc  must  be  ionized  every  hour  in  a  cell  in 
order  to  produce  a  current  of  5  amperes  strength?     For  how  long 
would  500  g.  of  zinc  serve  to  maintain  this  current? 

5.  What  will  be  the  E.M.F.  of  a  concentration  cell  in  which  the 
poles  are  of  lead  and  the  lead-ion  is  one  hundred  times  more  concen- 
trated round  one  pole  than  round  the  other? 

6.  What  weight  of  aluminium  must  become  ionized  every  hour 
in  a  cell  in  order  that  a  current  of  five  amperes  strength  may  be  pro- 
duced?    What  would  be  the  E.M.F.  of  the  current  if  an  acid  with 
normal  concentration  of  hydrogen-ion  surrounded  the  cathode  and  a 
solution   of   normal   aluminium-ion   the   anode?     How   would   this 
E.M.F.  be  affected  if  the  aluminium-ion  were  only  one-hundredth 
normal? 


CHAPTER  XXXIX 
ALUMINIUM  AND  THE  METALLIC  ELEMENTS  OF  THE  EARTHS 

THE  fourth  column  of  the  periodic  table  contains  boron  and 
aluminium  along  with  a  number  of  rare  elements.  The  chief 
members  of  the  family  are:  boron  (B,  at.  wt.  11),  aluminium  (Al, 
at.  wt.  27.1),  gallium  (Ga,  at.  wt.  69.9),  indium  (In,  at.  wt.  114.8), 
thallium  (Tl,  at.  wt.  204),  all  on  the  right  side  of  the  column;  and 
scandium  (Sc,  at.  wt.  44.1),  yttrium  (Yt,  at.  wt.  88.7),  lanthanum 
(La,  at.  wt.  139),  on  the  left  side.  These  elements  are  all  trivalent. 

The  Rare  Elements  of  this  Family.  —  The  oxide  and  hydrox- 
ide of  boron  are  acidic  (p.  637).  Those  of  aluminium  A1(OH)3,  gal- 
lium Ga(OH)3,  indium  In(OH)3,  and  thallium  T1O.OH  are  basic,  but 
behave  also  as  acids  towards  strong  bases. 

Gallium  and  indium  occur  occasionally  in  zinc-blende,  and  were 
discovered  by  the  use  of  the  spectroscope.  The  former  takes  its 
name  from  the  country  (France)  in  which  the  discovery  was  made, 
and  the  latter  from  two  blue  lines  shown  by  its  spectrum.  Indium 
gives  a  complete  series  of  compounds  in  which  it  is  trivalent,  and  the 
chlorides  InCl  and  InCl2  are  also  known. 

Thallium  is  found  in  some  specimens  of  pyrite  and  blende.  It  was 
discovered  by  Crookes,  by  means  of  the  spectroscope,  in  the  selenif- 
erous  deposit  from  the  flues  of  a  sulphuric  acid  factory.  It  received 
its  name  from  the  prominent  green  line  in  its  spectrum  (Gk.  0oAAos, 
a  green  twig) .  It  gives  two  complete  series  of  compounds.  In  those 
in  which  it  is  trivalent  (thallic  salts),  it  resembles  aluminium  (q.v.). 
Thus,  the  salts  of  this  series  are  more  or  less  hydrolyzed  by  water. 
Univalent  thallium  recalls  both  sodium  and  silver.  Thallous  hydrox- 
ide T1OH  is  soluble,  and  gives  a  strongly  alkaline  solution,  but  the 
chloride  is  insoluble  in  cold  water.  The  solutions  of  the  thallous  salts 
are  neutral.  The  metal  is  displaced  from  its  salts  by  zinc. 

Of  the  elements  on  the  left  side  of  the  column,  scandium,  whose 
existence  and  properties  were  predicted  by  Mendelejeff  (p.  465),  is  the 
best  known.  The  metals  of  the  rare  earths,  of  which  it  is  one,  are 
found  in  rare  minerals  such  as  euxenite,  gadolinite,  orthite,  and  mona- 

807 


808  INORGANIC  CHEMISTRY 

zite,  which  occur  in  Sweden,  Greenland,  and  the  United  States. 
Cerium  (Ce,  at.  wt.  140.25)  neodymium  (Nd,  at.  wt.  144.3),  and  pras- 
eodymium (Pr,  at.  wt.  140.9)  occur  along  with  lanthanum  in  cerite,  a 
silicate  of  these  four  elements.  These  four  are  included  amongst  the 
metals  of  the  rare  earths.  The  compounds  of  many  of  these  rare  ele- 
ments behave  so  much  alike  that  separation  is  difficult.  It  is  certain, 
however,  that  there  are  several  with  atomic  weights  near  to  that  of 
lanthanum  for  which  accommodation  cannot  easily  be  found  in  the 
periodic  table.  Ostwald  has  compared  them  to  a  group  of  minor 
planets  such  as,  in  the  solar  system,  takes  the  place  of  one  large 
planet. 

ALUMINIUM  AL 

The  Chemical  Relations  of  the  Element.  —  Aluminium  is 
trivalent  exclusively.  Its  hydroxide,  like  that  of  zinc  (p.  771),  is 
amphoteric,  that  is  to  say,  it  is  feebly  acidic  as  well  as  basic,  and  hence 
the  metal  forms  two  sets  of  compounds  of  the  types  Na3.AlO3  and 
A12.  (804)3.  The  salts  of  both  series  are  more  or  less  hydrolyzed  by 
water,  the  former  very  conspicuously  so.  It  is  worth  noting  that  the 
hydroxides  of  the  trivalent  metals,  or  metals  in  the  trivalent  con- 
dition, such  as  A1(OH)3,  Cr(OH)3,  Fe(OH)3,  are  all  distinctly  less 
basic  than  are  those  of  the  bivalent  metals  such  as  Zn(OH)2,  Cd(OH)2, 
Fe(OH)2,  Mn(OH)2.  This  fact  is  used  in  analysis  (cf.  also  p.  784)  in 
separating  the  two  sets.  When  a  solution  of  the  chlorides  is  shaken 
with  precipitated  barium  carbonate,  the  free  acid  from  the  more 
highly  hydrolyzed  salts  of  Al+++,  CT+++)  and  Fe"1"1"*  interacts  with  this 
substance,  the  hydrolysis  is  promoted: 

A1C13  +  3H20  <±  A1(OH)3  +  3HC1, 

and  eventually  the  hydroxides  A1(OH)3,  Cr(OH)3,  and  Fe(OH)3  are 
completely  precipitated.  The  chlorides  of  the  bivalent  metals  re- 
main in  the  solution.  Aluminium  does  not  enter  into  complex 
anions  or  cations.  In  this  it  differs  from  zinc  and  resembles  magne- 
sium. It  is  too  feebly  base-forming  to  give  salts  like  the  carbonate  or 
sulphite,  because  hydrolysis  causes  precipitation  of  the  hydroxide 
(p.  656,  see  p.  814). 

Occurrence.  —  Aluminium  is  found  very  plentifully  in  combina- 
tion, coming  next  to  oxygen  and  silicon  in  this  respect.  The  felspars 
(such  as  KAlSi3Os),  the  micas  (such  as  KAlSiOJ,  and  kaolin  (clay) 
H2Al2(Si04)2,H2O  are  the  commonest  minerals  containing  it.  Since 


ALUMINIUM  AND   METALLIC  ELEMENTS  809 

the  soil  has  been  formed  largely  by  the  weathering  of  minerals  like 
the  felspars,  clay  and  other  products  of  the  decomposition  of  such 
minerals  constitute  a  large  part  of  it.  Garnets,  which  are  found  in 
metamorphic  rocks,  are  mainly  an  orthosilicate  of  calcium  and 
aluminium  Ca3Al2(SiO4)3.  Turquoise  is  a  hydrated  phosphate 
A12(OH)3PO4,H20,  and  cryolite  a  double  fluoride  3NaF,AlF3.  Vari- 
ous forms  of  the  oxide  and  hydroxide  are  also  found. 

Preparation  and  Physical  Properties.  —  The  metal  is  now 
made  on  a  large  scale  by  electrolysis  of  the  oxide  A12O3  dissolved  in 
a  bath  of  molten  cryolite  (m.-p.  1000°),  a  process  invented  by  C.  M. 
Hall  (1886).  The  operation  is  conducted  in  cells  (5X3  feet,  or 
larger),  the  carbon  linings  of  which  form  the  cathodes  (Fig.  157). 
The  anodes  are  rods,  of  carbon,  which  + 
combine  with  the  oxygen  as  it  is  liber- 
ated. The  metal  (m.-p.  659°)  sinks  to 
the  bottom  of  the  cell  and  is  drawn  off 
periodically,  while  fresh  portions  of  the 
oxide  are  added  from  time  to  time.  The 
oxide  is  made  from  bauxite  (see  below), 
and  must  be  free  from  oxide  of  iron  and 
other  impurities,  as  the  metal  cannot  be 
purified  commercially.  A  current  den- 

V,         »  -  f        ^      i  FIG.  157. 

sity  of  5  amperes  per  sq.  cm.  of  cathode 

area  and  an  E.M.F.  of  5-6  volts  maintain  the  temperature  of  the 
molten  materials,  and  cause  the  decomposition.  In  1866  aluminium 
cost  $250-750  (£50-150)  per  kilogram.  In  1883  the  whole  produc- 
tion was  about  40  kilos.  In  1913  the  United  States  alone  consumed 
35  million  kilos,  costing  about  50  cents  (2/-)  per  kilo. 

The  metal  melts  at  658.5°,  but  is  not  mobile  enough  to  make 
castings.  It  is  exceedingly  light  (sp.  gr.  2.6),  and  in  tensile  strength 
excels  the  other  metals,  with  the  exception  of  iron  and  copper.  It  is 
malleable,  and  the  foil  is  taking  the  place  of  tin-foil  for  wrapping 
foods.  It  has  a  silvery  luster,  and  tarnishes  very  slightly,  the  tough, 
firmly  adhering  film  of  hydrated  oxide  first  formed  protecting  its  sur- 
face. Although,  comparing  cross-sections,  it  is  not  so  good  a  conduc- 
tor of  electricity  as  is  copper,  yet  weight  for  weight  it  conducts  better. 
It  is  difficult  to  work  on  the  lathe  or  to  polish,  because  it  sticks  to  the 
tools,  but  the  alloy  with  magnesium  (about  2  per  cent),  called 
magnalium,  has  admirable  qualities  in  these  respects.  Aluminium 
bronze  (5-12  per  cent  aluminium)  is  easily  fusible,  has  a  magnificent 


810  INORGANIC  CHEMISTRY 

golden  luster,  and  possesses  mechanical  and  chemical  resistance 
exceeding  that  of  any  other  bronze.  The  metal  and  its  alloys  are 
used  for  making  cameras,  opera-glasses,  cooking  utensils,  and  other 
articles  requiring  lightness  and  strength,  such  as  air-ships,  automobile 
and  bicycle  parts,  and  brewing  vats,  as  well  as  for  the  transmission  of 
electric  currents.  The  powdered  metal,  mixed  with  oil,  is  used  in 
making  a  silvery  paint. 

Chemical  Properties.  —  The  metal  displaces  hydrogen  from 
hydrochloric  acid  very  easily.  In  sulphuric  and  nitric  acid,  however, 
it  receives  a  coating  of  the  hydroxide,  formed  by  hydrolysis  of  the 
salt,  and  the  action  is  slow  in  the  former  case,  and  almost  nil  in 
the  latter.  It  displaces  hydrogen  also  from  boiling  solutions  of 
the  alkalies,  forming  aluminates:  !o  *>] 

2A1  +  6NaOH  *+  2Na3A103  +  3H2. 

In  consequence  of  its  very  great  affinity  for  oxygen,  aluminium 
displaces  from  their  oxides  all  the  metals  below  magnesium  in  the 
E.-M.  series.  Thus,  when  a  mixture  of  aluminium  powder  and 
ferric  oxide  (thermite)  is  placed  in  a  crucible  and  ignited  by  means 
of  a  piece  of  burning  magnesium  ribbon,  aluminium  oxide  and  iron 
are  formed: 

2A1  ->  A1203  +  2Fe. 


The  very  high  temperature  (about  3000°)  produced  by  the  action  is 
sufficient  to  melt  both  the  iron  (m.-p.  1530°)  and  the  oxide  of  alumin- 
ium (m.-p.  2050°).  The  products,  not  being  miscible,  separate  into 
two  layers.  This  very  simple  method  of  making  pure  specimens  of 
metafe  like  chromium,  uranium,  and  manganese,  whose  oxides  are 
otherwise  hard  to  reduce,  is  called  by  Goldschmidt,  the  inventor, 
aluminothermy.  By  preheating  the  ends  of  steel  rails  with  a  gasoline 
torch,  firing  a  mass  of  thermite  in  a  crucible  above  the  joint,  and  allow- 
ing the  iron  to  flow  into  the  joint,  perfect  welds  are  made.  In  the 
same  way,  large  castings,  like  propeller  shafts,  when  broken,  can  be 
mended.  The  sulphides,  such  as  pyrite,  are  reduced  with  equal 
vigor  by  aluminium. 

The  largest  part  of  the  aluminium  of  commerce  is  used  by  steel- 
makers. When  added  in  small  amount  (  <  1  :  1000)  to  molten  steel, 
it  combines  with  the  gases,  and  gives  sound  ingots  free  from  blow 
holes. 


ALUMINIUM   AND  METALLIC  ELEMENTS  811 

Aluminium  Chloride  AlCk.  —  If  the  metal  or  the  hydroxide 
is  treated  with  hydrochloric  acid,  and  the  solution  is  allowed  to  evap- 
orate, crystals  of  A1C13,6H2O  are  formed.  When  heated,  this  hydrate 
is  completely  hydrolyzed,  hydrochloric  acid  is  given  off,  and  only  the 
oxide  remains.  The  anhydrous  chloride  is  much  used  as  a  catalytic 
agent  for  causing  combination  in  organic  chemistry.  It  is  made  by 
passing  dry  chlorine  over  aluminium. 

Aluminium  chloride  gives  a  vapor  pressure  of  760  mm.  at  183°, 
and  sublimes,  as  a  white  crystalline  solid,  without  melting.  Under 
pressure,  it  melts  at  193°.  In  the  mode  of  preparation  described 
above,  it  is,  therefore,  vaporized,  and  condenses  in  a  cool  part  of  the 
tube.  It  fumes  when  exposed  to  moist  air,  on  account  of  the  hydro- 
gen chloride  produced  by  hydrolysis,  and  only  with  excess  of  hydro- 
chloric acid  does  it  give  a  clear  solution  free  from  basic  salts. 

Aluminium  Hydroxide  and  the  Aluminates.  —  When  an 
alkali  is  added  to  a  solution  of  a  salt  of  aluminium,  the  hydroxide 
A1(OH)3  is  precipitated  in  gelatinous  form.  It  loses  water  gradually 
when  dried,  without  forming  any  intermediate  hydroxides  (p.  634), 
until  A12O3  alone  remains.  Natural  forms  of  this  substance  are 
hydrargyllite  A1(OH)3  (=  A12O3,3H2O),  bauxite  A120(OH)4  (=  A12O3, 
2H20),  which  always  contains  ferric  oxide,  and  diaspore  A10.OH 
(=  A1203,H20). 

Commercially,  the  hydroxide  is  made  by  heating  bauxite  with 
dry  sodium  carbonate,  or  with  concentrated  sodium  hydroxide  solu- 
tion at  150-170°.  The  ferric  oxide,  having  no  tendency  to  form  a 
carbonate  or  to  interact  with  a  base,  remains  unchanged.  The 
sodium  aluminate  which  is  formed  is  extracted  with  water: 

A12O(OH)4  +  Na2CO3  -»  2NaAlO2  +  CO2  +  2H2O. 

The  hydroxide  is  then  precipitated  by  passing  carbon  dioxide  through 
the  solution: 

2NaA102  +  CO2  +  3H20  ->  Na*CO3  +  2A1(OH)3. 

Aluminium  hydroxide,  being  amphoteric,  interacts  both  with 
acids  and  with  bases,  and  is,  therefore,  like  zinc  hydroxide  (p.  771), 
ionized  both  as  a  base  and  as  an  acid.  It  interacts  only  slightly 
with  ammonium  hydroxide,  because  this  substance  is  too  feebly 
basic,  but,  from  the  solution  in  the  active  alkalies,  the  aluminates 


812  INORGANIC  CHEMISTRY 

Na3.A103,  Na.A102,  and  K.A102  can  be  obtained  in  solid  form. 
The  aluminates  are  largely  hydrolyzed  by  water: 

NaA102  +  2H20  <=±  NaOH  +  A1(OH)3. 

Hence  an  excess  of  sodium  hydroxide  is  required  for  the  complete 
solution  of  aluminium  hydroxide  by  the  reversal  of  this  action.  So- 
dium aluminate  is  used  as  a  mordant  in  dyeing  (see  below),  on  ac- 
count of  the  ease  with  which  the  solution  gives  up  aluminium 
hydroxide  when  any  material  is  present  which  can  adsorb  the  free 
portion  of  the  hydroxide  and  so  cause  forward  displacement  of  the 
above  equilibrium. 

When  calcium  chloride  is  added  to  a  solution  of  sodium  aluminate, 
the  insoluble  calcium  metaluminate  is  deposited: 

2NaAlO2  +  CaCl2  -»  Ca(AlO2)2  +  2NaCl. 
The  relations  of  these  substances  are  shown  by  the  following  formulae: 


,0-H          ,0-Na  Q  0  X 

Al-O-H    Al-O-Na    Al  '  Al  (  .Ca. 

^0-H         X0-Na  °-H  °-Na        /° 


A  number  of  insoluble  metaluminates  are  found  in  nature.  They 
crystallize  in  the  regular  system,  and  are  known  as  spinelles.  They 
contain  bivalent  metals  in  place  of  the  calcium  in  the  last-named 
compound.  Thus  we  have  spinelle  proper  Mg(AlO2)2,  and  gahnite 
Zn(AlO2)2.  Corresponding  and  isomorphous  derivatives  of  chromic 
and  ferric  hydroxides  are  chromite  Fe(Cr02)2  and  magnetite  Fe 
(Fe02)2. 

Aluminium  Oxide  Al^O*.  —  The  oxide  (alumina)  is  made  by 
heating  the  pure  hydroxide  made  from  bauxite  (see  above).  It  is 
found  in  nature  in  pure  form  as  corundum.  This  mineral  is  only 
one  degree  less  hard  than  the  diamond.  Emery  is  a  common  variety, 
contaminated  with  ferric  oxide,  and  was  widely  used  as  an  abrasive 
until  largely  displaced  by  carborundum.  The  ruby  is  pure  aluminium 
oxide  tinted  by  a  trace  of  a  compound  of  chromium,  while  the  sap- 
phire is  the  same  material  colored  with  aluminates  of  iron  and  tita- 
nium. It  is  said,  however,  that  the  same  tint  is  conferred  upon 


ALUMINIUM  AND  METALLIC  ELEMENTS  813 

colorless  corundum  by  exposure  to  the  influence  of  salts  of  radium. 
By  allowing  the  pulverized  oxides  to  be  carried  by  the  stream  of 
oxygen  of  an  oxy-hydrogen  blowpipe,  and  placing  a  disc  in  the  flame 
to  catch  the  molten  particles,  "synthetic"  sapphires  and  rubies  are 
now  made  in  large  quantities.  Alundum,  a  refractory  material  for 
crucibles,  is  made  by  heating  objects  made  of  the  oxide  in  the  electric 
furnace  until  a  small  proportion  of  the  material  is  melted.  The 
alumina  made  by  gently  heating  the  hydroxide  interacts  easily  with 
acids,  but  after  being  strongly  heated  it  resembles  natural  alumina 
in  being  very  slowly  affected  by  them.  Minerals  containing  insoluble 
compounds  of  aluminium  are  attacked  when  heated  strongly  with 
potassium  bisulphate  (cf.  p.  674),  the  sulphate  of  aluminium  being 
formed. 

Aluminium  Sulphate:    The  Alums.  —  Aluminium  sulphate 

A12(SO4)3,18H2O  is  prepared  by  treating  either  bauxite  or  pure  clay 
(kaolin)  with  sulphuric  acid.  In  the  latter  case  the  insoluble  resi- 
due of  silicic  acid  is  removed  by  filtration: 


+  3H2S04  ->  A12(S04)3  +  2H2Si03  +  2H20. 

The  solution  of  the  sulphate  is  acid  in  reaction.  It  crystallizes 
in  leaflets  which,  when  the  source  was  clay  or  bauxite,  have  a  yellow 
tinge  due  to  the  presence  of  iron  as  an  impurity.  The  salt  is  used 
as  a  source  of  precipitated  aluminium  hydroxide  in  paper-making, 
water  purification,  and  dyeing. 

When  potassium  sulphate  solution  is  added  to  a  strong  solution 
of  aluminium  sulphate,  octahedral  crystals  of  potash  alum,  K2SO4, 
A12(SO4)3}24H2O  are  deposited.  This  is  a  double  salt,  and  is  one  of 
a  large  number  known  as  the  alums.  These  have  the  general  formula 
M2ISO4,M2III(SO4)3,24H20,  and  may  be  made  as  above  by  using  a 
sulphate  of  a  univalent  metal  with  one  of  a  trivalent  metal.  Thus, 
for  M1  we  may  use  K,  NH4,  Rb,  Cs,  and  Tl1,  and  for  Mm,  Al, 
Fem,  Cr111,  Mnm,  and  Tlm.  We  may  even  employ  selenates,  such 
as  K2SeO4.  All  of  the  resulting  double  salts  are  isomorphous,  and 
a  crystal  of  one  will  continue  to  grow  in  a  solution  of  another, 
acquiring,  of  course,  an  outer  layer  of  different  composition,  but  of 
the  same  crystallographic  orientation. 

Potassium-  Aluminium  Sulphate  KzSOi9Al^(SO^)3924HzO.  —  • 
Ordinary  alum  is  made  from  aluminium  sulphate  obtained  from 
clay  (see  above).  It  is  also  prepared  by  heating  alunite,  a  basic 


814  INORGANIC  CHEMISTRY 

alum  found  near  Rome  and  in  Hungary,  and  extracting  the  product 
with  hot  water.  The  alunite,  having  the  composition  KA13(OH)6 
(804)2,  leaves  an  insoluble  residue  of  the  hydroxide,  mixed  with  ferric 
oxide  which  is  present  as  an  impurity: 

2KA13(OH)6(SO4)2  -*  K2S04,A12(S04)3  +  4A1(OH)8. 

The  saturated  solution  of  alum  contains,  at  10°,  9  parts  of  the 
anhydrous  salt  in  100  parts  of  water,  and  at  100°,  422  parts  in  100  of 
water.  The  hydrated  salt  melts  at  90°.  An  aqueous  solution  of  this 
salt  or  of  sodium  phosphate  (p.  559)  is  used  for  fire-proofing  draperies, 
because  the  crystals  deposited  in  the  fabric  melt  easily,  and  the  fused 
material  protects  the  fibers  from  access  of  oxygen.  When  heated 
more  strongly  alum  loses  its  water  of  hydration,  together  with 
some  sulphur  trioxide,  and  leaves  a  slightly  basic,  anhydrous  salt 
known  as  burnt  alum.  A  solution  of  alum  dissolves  a  considerable 
amount  of  aluminium  hydroxide,  giving  " neutral  alum,"  a  basic  salt 
K2SO4,Al4(OH)6(SO4)3  used  as  a  mordant  (see  p.  820).  The  substance 
is  usually  prepared  by  adding  sodium  carbonate  to  the  solution  of 
alum  as  long  as  the  aluminium  hydroxide,  formed  locally,  continues 
to  redissolve.  Potash-alum  and  ammonium-alum  are  more  easily 
freed  from  impurities  (e.g.,  compounds  of  iron)  by  recrystallization 
than  is  aluminium  sulphate,  and  the  alums  are  therefore  used  instead 
of  the  latter  in  medicine,  in  dyeing  delicate  shades,  and  to  replace 
cream  of  tartar  in  baking  powder  (p.  689).  In  the  last  case,  the 
reaction 

K2SO4,A12(S04)3,24H20  +  6NaHC03^ 
K2SO4  +  3Na2SO4  +  2A1(OH)3  +  6CO2  +  24H2O 

liberates  carbon  dioxide  by  hydrolysis  of  the  aluminium  carbonate. 

Hydrolysis  of  Aluminium  Carbonate.  —  The  foregoing  ac- 
tion, and  others  discussed  above  (p.  808),  show  that  the  carbonate  is 
completely  hydrolyzed: 

A12(C03)3  +  6H20  <F±  2A1(OH)3  J  +  3H2C03  -*  3H20  +  3C02  T . 

It  will  be  seen  that  this  may  be  due  only  in  part  to  the  feebly  basic 
qualities  of  the  hydroxide.  If  the  hydroxide  were  not  precipitated, 
it  would  cause  some  reversal  of  the  action,  and  some  of  the  car- 
bonate would  remain.  The  insolubility  of  one  product  explains 
also  other  cases  of  the  complete  hydrolysis  of  salts  (e.g.,  ammonium 
silicate,  p.  648,  and  next  section). 


ALUMINIUM  AND   METALLIC  ELEMENTS  815 

Aluminium  Sulphide  AlA.  —  This  compound  is  most  easily 
obtained  by  mixing  pyrite  with  aluminium  powder  and  igniting  with 
magnesium  ribbon  (p.  653) :  3FeS2  +  4A1  ->  2A12S3  +  3Fe.  It  forms 
a  grayish-black  solid,  and  is  hydrolyzed  by  water  as  is  magnesium 
sulphide,  giving  the  hydroxide  and  hydrogen  sulphide.  On  this 
account,  the  sulphide,  like  magnesium  sulphide  (p.  767),  cannot  be 
formed  by  precipitation  in  presence  of  water.  Thus,  ammonium 
sulphide  with  a  salt  of  aluminium,  in  solution,  gives  a  precipitate 
of  aluminium  hydroxide: 

A12(SO4)3  +  3(NH4)2S  +  6H20  ->  2A1(OH)3 1  +  3(NH4)2SO4  +  3H2S. 

Aluminium  Acetate.  —  This  salt  is  used  by  dyers,  because, 
being  a  salt  of  a  weak  base  and  a  weak  acid,  it  is  much  hydrolyzed  by 
water,  especially  at  100°.  In  mordanting,  it  thus  gives  aluminium 
hydroxide  very  easily.  It  is  made  by  treating  lead  or  barium  acetate 
with  aluminium  sulphate,  and  filtering  and  crystallizing  the  solution: 

A12(SO4)3  +  3Ba(CO2CH3)2^3BaSO4 j  +  2A1(CO2CH3)3. 

Coagulation  Method  of  Purifying  Water:   Sizing  Paper.  — 

When  aluminium  hydroxide  is  formed,  it  gives  first  a  colloidal  sus- 
pension, which  coagulates  to  a  gelatinous  precipitate.  When  this 
precipitate  is  produced  in  water  used  for  domestic  purposes,  and 
containing  fine,  suspended  matter,  the  gelatinous  material  causes 
the  fine  particles  to  collect  into  larger  ones  which  settle  rapidly, 
and  permits  the  use  of  relatively  small  settling  ponds.  These 
larger  particles  enclose  also  practically  all  the  bacteria.  If  the 
water  is  slightly  hard,  crude  aluminium  sulphate,  alone,  is  used: 

3Ca(HC03)2  +  A12(SO4)3  -*  3CaS04  +  2A1(HCO3)3          (1) 
A1(HC03)3  +  3H20  ->  A1(OH)8  |  +  3H2CO3  (2) 

If  the  water  is  very  soft,  a  little  lime  Ca(OH)2  is  added.  The  few 
remaining  organisms  (particularly  colon  bacilli,  p.  142)  are  destroyed 
by  later  addition  of  bleaching  powder  or  of  chlorine-water  (p.  226). 

In  some  localities  crude  ferrous  sulphate,  obtained  as  a  by-product 
in  cleaning  iron,  is  cheaper,  and  is  employed  instead.  The  lime 
precipitates  ferrous  hydroxide  Fe(OH)2.  This  is  quickly  oxidized 
to  colloidal  ferric  hydroxide  Fe(OH)3,  which  coagulates  the  suspended 
matter. 

Aluminium  hydroxide  is  employed  also  in  sizing  cheaper  grades 
of  paper  (p.  603),  an  operation  required  to  prevent  the  absorption 


816  INORGANIC  CHEMISTRY 

and  consequent  spreading  of  the  ink.  For  writing-paper,  gelatine 
solution  is  employed.  In  making  printing-papers,  rosin  soap  (made 
by  dissolving  rosin  in  caustic  soda)  is  mixed  with  the  paper-pulp, 
and  aluminium  sulphate  is  added.  The  rosin  and  aluminium  hydrox- 
ide are  precipitated  in  the  pulp,  and  pressing  between  hot  rollers 
afterwards  melts  the  former  and  gives  a  surface  to  the  paper. 

Delicate  cloth  goods  are  rendered  waterproof  by  saturating  them 
with  aluminium  acetate  solution  and  then  steaming  them  to  promote 
hydrolysis.  The  aluminium  hydroxide  is  thus  precipitated  in  the 
capillaries  of  the  cotton  or  linen  and  renders  them  non-absorbent: 

A1(C02CH3)3  +  3H20  <=»  A1(OH)8  +  3HC02CH3. 

Kaolin  and  Clay:    Earthenware  and  Porcelain.  —  By  the 

action  of  water  and  carbon  dioxide  upon  granite,  and  other  rocks 
containing  felspar  KAlSi3Os,  the  potash  is  slowly  removed,  and 
the  compound  is  changed  largely  into  a  hydrated  orthosilicate 
H2Al2(Si04)2,H2O.  When  pure,  it  forms  kaolin  or  china  clay,  a 
white,  crumbly  material.  When  washed  away  and  redeposited,  it 
usually  acquires  compounds  of  iron,  the  carbonates  of  calcium  and 
magnesium,  and  sand  (silica),  becoming  common  clay.  Ocher, 
umber,  and  sienna  are  clays  colored  with  oxides  of  iron  and  man- 
ganese. Fuller's  earth  is  a  purer  variety. 

The  plasticity  of  clay,  a  property  connected  with  the  colloidal 
nature  of  the  kaolin,  enables  it  to  be  fashioned  into  various  shapes. 
When  heated,  it  shrinks  and  becomes  a  hard,  solid,  porous  mass, 
and  does  not  melt.  These  two  properties  enable  us  to  use  it  in 
making  bricks,  pottery,  and  porcelain.  The  presence  of  calcium  and 
magnesium  compounds  makes  the  clay  more  fusible,  because  it 
permits  the  formation  of  fusible  silicates  of  these  metals.  Bricks 
and  tiling  for  roofs  and  drains  are  made  of  common  clay  and,  when 
red,  owe  their  color  to  oxide  of  iron  Fe203.  The  firing  is  done  with 
fuel  gas  in  ovens  or  kilns  of  brickwork.  The  efflorescence  which  often 
appears  on  the  surface  of  the  bricks  (" niter")  is  generally  due  to 
sodium  sulphate  or  sodium  chloride  present  originally  in  the  clay. 
To  glaze  drain  pipes  and  some  bricks,  salt  is  thrown  into  the  kiln. 
The  water  vapor,  at  a  red  heat,  hydrolyzes  the  vapor  of  the  salt, 
hydrogen  chloride  is  set  free,  and  the  sodium  hydroxide  gives  with 
the  clay  a  fusible  sodium-aluminium  silicate  which  fills  the  surface 
pores.  Clay  for  fire  brick  (infusible)  must  contain  silica,  but  no 
lime. 


ALUMINIUM   AND  METALLIC  ELEMENTS  817 

China  and  porcelain  are  made  from  pure  clay,  free  from  iron, 
to  which  a  little  of  the  more  fusible  felspar  is  added.  After  the 
first  firing,  the  articles  are  porous  (bisque),  and  must  be  covered 
with  a  glaze  to  make  them  water-tight.  A  paste  of  finely  ground 
felspar  and  silica,  sometimes  containing  lead  oxide,  is  spread  on  the 
surface,  and  the  articles  are  fired  again,  at  a  higher  temperature. 
The  felspar  melts  and  fills  the  pores,  so  that  a  continuous,  semi- 
transparent  material  results.  Colored  decorations  are  added  by 
means  of  suitable  materials,  mainly  oxides  of  metals  which  give 
colored  silicates.  The  third  firing  causes  these  oxides  to  interact 
and  fuse  with  the  glaze.  Porcelain,  if  made  with  sufficient  clay,  is 
very  infusible.  It  is  attacked  by  aqueous  and  by  fused  alkalies, 
however,  giving  soluble  silicates. 

The  Schwerin  process  for  separating  ferric  oxide  Fe2Os  from 
clay,  so  that  white  porcelain  may  be  obtained,  is  now  used  on  a 
large  scale  and  affords  an  interesting  application  of  the  properties 
of  colloidal  suspensions  (p.  622).  When  the  impure  clay  is  sus- 
pended in  water,  the  particles  of  ferric  oxide  are  positively  charged 
and  those  of  the  clay  are  negative.  By  inserting  plates  connected 
with  the  dynamo  in  the  trough,  the  clay  particles  are  caused  to 
drift  towards  the  positive  plate  and  the  ferric  oxide  towards  the 
other,  so  that,  when  the  liquid  from  the  positive  end  is  allowed  to 
settle,  pure  clay  is  obtained. 

In  making  bricks,  in  some  cases,  advantage  is  taken  of  the  fact 
that  negative  colloids,  such  as  clay,  become  more  strongly  negative 
in  presence  of  a  trace  of  free  alkali.  Thus,  when  a  trace  of  sodium 
hydroxide  is  added  to  clay  slip,  the  particles  repel  one  another  more 
strongly,  the  cohesion  which  causes  the  plasticity  is  reduced,  and 
the  clay  can  be  poured  into  molds.  This  avoids  diluting  the  clay 
with  water,  which  would  only  have  to  be  driven  out  again,  with 
great  waste  of  heat,  in  the  firing. 

Cement.  —  Cement  is  made  by  heating  limestone  CaCOs,  and 
clay  HAlSi04,  or  a  natural  rock  containing  both  materials  in 
the  right  proportions.  Such  a  rock,  made  into  cement  by  volcanic 
heat,  was  quarried  by  the  Romans  near  Naples  and  elsewhere,  and 
its  capacity  for  hardening  even  under  water  was  utilized  by  them. 
Blast-furnace  slag,  when  pulverized  and  heated  with  limestone,  has 
been  found  to  yield  an  excellent  quality  of  cement,  and  a  valuable 
use  has  thus  been  found  for  what  was  formerly  an  annoying  encum- 
brance. The  mixture,  or  the  pulverized  natural  rock,  is  moistened 


818 


INORGANIC  CHEMISTRY 


and  fed  slowly  in  at  the  upper  end  of  an  inclined  (6°)  revolving 
cylinder  of  iron  (20  to  45  by  2  meters).  The  motion  continually 
turns  over  the  thin  layer,  and  exposes  every  particle  to  the  heat  of 
the  air-blast,  charged  with  pulverized  coal,  burning  in  the  interior. 
The  product  slides  out  in  a  continuous  stream  at  the  lower  end,  and  is 
pulverized  by  steel  balls  in  a  ball  mill. 

Cement  is  held  to  be  a  mixture  of  calcium  silicate  and  calcium 
aluminate.  The  former  is  simply  a  filler.  The  latter  is  hydrolyzed 
by  the  water: 

Ca3(A103)2  +  6H20  ->  3Ca(OH)2 


The  calcium  hydroxide  slowly  crystallizes,  connecting  the  particles 
of  the  calcium  silicate.  The  aluminium  hydroxide  fills  the  inter- 
stices and  renders  the  whole  compact  and  impervious. 

Ultramarine.  —  Formerly,  pulverized  lapis  lazuli,  a  rare  min- 
eral of  beautiful  blue  color,  was  used  by  artists  as  a  pigment.  Gmelin 
(1828)  found  a  way  of  making  it  artificially.  A  mixture  of  kaolin, 
sodium  carbonate,  sulphur,  and  charcoal  is  heated  until  a  green  mass 
is  obtained.  The  mass  is  then  pulverized  and  heated  with  more 
sulphur.  Its  composition  is  approximately  4NaAlSi04,Na2S2.  The 
product  is  used  as  laundry  blueing,  in  making  blue-tinted  paper,  and 
with  oil  in  making  paint.  It  is  also  added  in  small  amount  to  correct 
the  natural  yellow  tint  of  linen,  starch,  sugar  (p.  606),  and  paper-stock. 
By  varying  the  mode  of  heating,  without  altering  the  composition, 
various  colors  from  green  to  reddish  violet  can  be  obtained.  No 
pure  colored  substance  can  be  extracted  from  it.  The  variety  of 
colors  is  due  to  different  degrees  of  colloidal  dispersion  of  some  sub- 

stance suspended  in  the  solid,  just 
as  gold,  which  is  pale  yellow  in 
mass,  gives  colloidal  suspension  (p. 
621)  of  different  colors  (red,  purple, 
or  blue)  according  to  the  fineness 
of  the  particles  (cf.  p.  750). 


Dyeing.  —  The  problem  of  the 
dyer  is  to  confer  the  desired  color 
FlG  158  upon  a  fabric  made,  usually,  of  cot- 

ton, linen,  wool,  or  silk,  and  to  do 

this  in  such  a  way  that  the  dye  is  fast  to  (i.e.,  is  not  removed  or 
destroyed  by)  rubbing  and  light,  and  often,  also,  to  washing  with 


ALUMINIUM   AND  METALLIC  ELEMENTS  819 

soap.  To  understand  the  means  by  which  this  is  achieved,  it  must 
be  noted  that  cotton  and  linen  consist  of  smooth  hollow  fibers  (Fig. 
158A)  of  the  composition  of  cellulose  (CeHioOs).,;.  Wool  is  made 
of  hollow  fibers,  with  a  scaly  surface  (B)  and  silk  (CisH^NsOe);,;  of 
solid  filaments,  but  both  are  composed  of  proteins  (p.  628).  Now, 
the  proteins  are  much  more  active  chemically  than  is  cellulose,  and 
also,  as  colloidal  materials,  seem  to  have  a  much  greater  tendency 
to  adsorb  other  substances  than  has  cellulose.  Hence,  accidental 
stains  on  wool  or  silk  are  much  less  often  removable  than  are  those 
on  cotton,  and  when  samples  of  the  three  materials  are  dipped  in  a 
solution  of  a  dye,  the  first  two  are  permanently  dyed,  while  from  the 
last  most  dyes  can  be  completely  washed  out  with  water. 

Three  modes  of  dyeing  may  be  mentioned  : 

1.  Insoluble  Dyes.  If  the  colored  body  can  be  produced  by 
precipitation,  after  the  solution  has  filled  the  capillary  and  wall 
of  every  fiber  of  the  goods,  then,  if  the  dye  is  sufficiently  insoluble, 
it  is  mechanically  imprisoned  in  every  fiber  and  cannot  be  washed 
out.  This  plan  may  be  applied  to  any  kind  of  goods.  For  example, 
if  cotton,  silk,  or  wool  is  first  boiled  in  a  solution  of  lead  acetate, 
and  is  then  soaked  in  a  boiling  solution  of  potassium  chromate, 
it  is  dyed  a  brilliant,  permanent  yellow.  Lead  chromate  is  the  col- 
ored body  : 


Pb(C02CH3)2  +  K2CrO4±+  2K(C02CH3)  +  PbCr04  1. 

The  part  precipitated  on  the  outside  of  the  goods  can  be,  and  is, 
at  once  washed  off  by  rubbing  in  water,  but  the  particles  inside  the 
fibers  can  come  out  only  by  being  dissolved,  and  they  are  insoluble 
in  watei.  Indigo  Ci6H]0N202,  which  is  used  in  larger  amounts  than 
any  other  dye,  belongs  to  this  class.  Obtained  in  early  times  from 
several  points  in  Europe  and  Egypt,  where  it  was  known  as  woad, 
and  more  recently  imported  from  India,  where  the  cultivation  of 
the  indigo  plant  was  as  important  an  industry  as  is  the  growing 
of  cotton  in  the  Southern  States,  it  is  now  almost  all  made  artificially. 
Synthetic  indigo  has  been  manufactured  since  1907,  with  naphtha- 
lene CioHg  (p.  613),  obtained  from  gas  tar  and  the  tar  from  by-product 
coke  ovens,  as  the  initial  substance.  The  cloth  is  saturated  with 
an  alkaline  solution  of  indigo  white  Ci6Hi2N2O2  (cf.  p.  444),  a  soluble, 
slightly  acid  substance,  and  the  oxygen  of  the  air  subsequently  oxi- 
dizes this  and  deposits  the  insoluble  indigo  blue  within  the  fibers  : 

2Ci6H12N202  +  02  ->  2Ci6H10N202i  +  2H2O. 


820  INORGANIC   CHEMISTRY 

Indanthrene  blue  is  applied  in  the  same  way  as  indigo,  and  is  even 
less  affected  by  light. 

2.  Mordant  or  Adjective  Dyes.     Since  cotton  is  inactive  chemi- 
cally and,  although  a  colloid,  has  but  a  slight  tendency  to  adsorb 
dyes,  it  is  usually  necessary  first  to  introduce  into  the  fibers  of  cot- 
ton some  colloidal  substance  with  greater  adsorptive  powers.     Sub- 
stances of  this  kind  are  tannic  acid  for  basic  dyes,  and  gelatinous 
colloidal  hydroxides,   such  as  those  of  aluminium,  tin,  iron,  and 
chromium,  for  non-basic   (including  acid)   dyes.     They  are  called 
mordants  (Lat.  mordere,  to  bite).     Thus,  if  in  three  jars  we  place 
very^dilute  solutions  of  aluminium  sulphate,  ferric  chloride  FeCl3 
and  chromous  acetate  Cr(C02CH3)2,  then  add  a  few  drops  of  a 
solution  of  a  dye  to  each,  and  finally  introduce  a  little  of  a  base 
(like  sodium  hydroxide)  to  precipitate  the  hydroxide  of  the  metal, 
this  hydroxide  will  adsorb  the  dye  and  carry  it  into  the  precipitate. 
Such  a  precipitate  of  mordant  and  dye  is  called  a  lake  (Fr.  laque,  lac) . 
With  the  same  dye,  the  three  lakes  have  different  colors.    Thus,  in 
the  above  mentioned  experiment,  if  alizarin  CnHgCU  (madder),  an 
orange-yellow,  slightly  soluble  acid,  is  used  as  the  dye,  the  colors 
are  red  (Turkey  red),  violet,  and  maroon,  respectively.     This,  of 
course,  is  due  to  the  different  degrees  of  dispersion  in  the  three 
colloidal  materials.     If  aluminium  hydroxide  is  to  be  used,  by  first 
saturating  the  cloth  with  hot  aluminium  acetate  solution  (p.  815), 
or  by  using  first  aluminium  sulphate  and  then  ammonium  hydroxide, 
the  aluminium  hydroxide  is  precipitated  within  the  fibers  of  the  goods. 
When  the  material  is  then  dyed,  the  coloring  matter  is  adsorbed  by 
the  mordant,  with  which  it  forms  an  insoluble  lake,  within  the 
fibers.     Cochineal,   obtained  from  wingless,   female   insects   found 
on  the  variety  of  cactus  bearing  this  name,  contains  carminic  acid 
CnHi207,  and  gives  a  red  lake  with  aluminium  hydroxide.     Basic 
dyes,  like  Malachite  green  and  Methylene  blue,  behave  similarly 
with  tannic  acid,  or  an  insoluble  salt  of  tannic  acid,  as  mordant. 
It  will  be  seen  that,  so  far  as  the  fabric  is  concerned,  this  process, 
like  the  first,  is  a  mechanical  one,  and  is  independent  of  the  chemical 
nature  of  the  goods. 

3.  Direct  or  Substantive  Dyes.     Most  organic  dyes  are  direct 
dyes  on  silk  or  wool,  and  require  no  mordant  with  these  materials. 
The  actions  seem  to  be  sometimes  chemical,  but  more  often  cases 
of  adsorption  by  the  silk  or  wool  (both  colloids)  themselves.     A 
few  dyes  are  also  fast  on  cotton.     Congo-red   NasCsuH^NeS-jOe  is 
fast  both  on  cotton  and  wool,  but  is  no  longer  much  used.     Chry- 


ALUMINIUM   AND  METALLIC  ELEMENTS  821 

Aophenin  is  now  one  of  the  commonest  dyes  of  this  class.  These 
dyes,  which  are  sodium  salts  of  complex  organic  acids,  are  col- 
loids like  soap  (p.  623)  and  are  salted  out  within  the  fibers  of  the 
goods  by  adding  sodium  sulphate  to  coagulate  them  and  assist  the 
adsorption  by  the  cotton.  Once  adsorbed  in  this  way,  unlike  soap, 
they  cannot  be  washed  out. 

Analytical   Reactions    of  Aluminium    Compounds.  —  The 

alkalies,  and  alkaline  solutions  like  that  of  ammonium  sulphide, 
precipitate  the  white  hydroxide  (p.  815).  The  product  is  soluble  in 
excess  of  the  active  alkalies.  Soluble  carbonates  also  throw  down 
the  hydroxide.  Aluminium  compounds,  when  heated  strongly  in 
the  flame  with  cobalt  salts,  give  a  blue  aluminate  of  cobalt  Co(A102)2, 
Thenard's  blue. 

Exercises.  —  1.   What  are  the  differences  between  zinc  and  alu- 
minium, and  their  corresponding  compounds? 

2.  Construct  equations  showing,  (a)  the  hydrolysis  of  aluminium 
sulphate  (p.  813),   (6)  the  interaction  of  aluminium  sulphate  and 
ecbslt  nitrate  in  the  Bunsen  flame. 

3.  Formulate  the  ionization  of  aluminium  hydroxide  (pp.  771, 
811). 

4.  Why  does  zinc  hydroxide,  in  spite  of  its  feebleness  as  a  base, 
dissolve  in  ammonium  hydroxide,  while  aluminium  hydroxide  does 
not? 


CHAPTER  XL 
GERMANIUM,   TIN,   LEAD 

THE  elements  on  the  right  side  of  the  fifth  column  of  the  periodic 
table,  aside  from  carbon  and  silicon,  are  germanium  (Ge,  at.  wt. 
72.5),  tin  (Sn,  at.  wt.  118.7),  and  lead  (Pb,  at.  wt.  207.2).  Titanium 
(Ti,  at.  wt.  48.1),  zirconium  (Zr,  at.  wt.  90.6),  cerium  (Ce,  at.  wt. 
140.25),  and  thorium  (Th,  at.  wt.  232.4)  occupy  the  left  side. 

The  Chemical  Relations  of  the  Family.  —  All  of  these  ele- 
ments show  a  maximum  valence  of  four.  Germanium,  tin,  and  lead 
are  also  bivalent.  In  this  respect  they  resemble  carbon  and  differ 
from  silicon,  which  is  more  closely  allied  to  the  elements  on  the  left 
side  of  the  column.  The  oxides  and  hydroxides  in  which  these  three 
elements  are  bivalent  become  more  basic,  and  the  elements  themselves 
more  metallic  in  chemical  relations,  with  increase  in  atomic  weight. 
In  this  they  resemble  the  potassium  and  calcium  families.  Curiously 
enough,  the  same  three  hydroxides  are  also  acidic.  They  are  more 
strongly  acidic  than  is  zinc  hydroxide,  for  the  salts  they  form  by 
interaction  with  bases  are  less  hydrolyzed  than  are  the  zincates. 
This  acidic  character  likewise  increases  in  the  order  in  which  the  ele- 
.ments  are  named  above. 

GERMANIUM  GE 

Germanium  (cf.  p.  465)  may  be  described  as  a  transition  element 
between  carbon  and  tin.  It  forms  two  oxides  GeO  and  Ge02  corre- 
sponding to  those  of  carbon  and  of  tin.  Germanious  oxide  is  not  very 
definitely  basic  or  acidic,  and  the  sulphide  is  the  only  other  well- 
defined  compound  of  this  set.  Germanic  oxide  and  hydroxide  are 
acidic  entirely.  The  resemblance  to  carbon  is  shown  in  the  forma- 
tion of  an  unstable  compound  with  hydrogen,  and  of  germanium 
chloroform  GeHCl3.  Like  carbon,  tin,  and  silicon,  germanium  gives 
a  volatile  chloride  GeCU  (b.-p.  87°).  Like  tin  and  gold  (p.  760), 
it  forms  complex  sulphides  derived  from  germanic  sulphide,  such 
as  K2GeS3.  The  element  was  discovered  (in  1886)  in  argyrodite,  a 
complex  sulphide  4Ag2S,GeS2. 

822 


GERMANIUM,   TIN,   LEAD  823 

TIN  SN 

The  Chemical  Relations  of  the  Element.  —  Tin  is  both  biva- 
lent and  quadrivalent.  Each  of  the  hydroxides,  Sn(OH)2  and 
SnO(OH)2  (or  Sn(OH)4),  is  both  basic  and  acidic,  so  that  there  are 
really  four  series  jof  compounds.  Still,  stannous  hydroxide  is  mainly 
a  base,  of  a  feeble  sort,  while  stannic  hydroxide  is  mainly  an  acid. 
Thus  we  have  stannous  chloride,  sulphate,  and  nitrate,  which  are 
stable,  although  they  are  all  more  or  less  hydrolyzed  by  water,  and 
sodium  stannite  Na2.SnO2  which  is  unstable.  On  the  other  hand, 
stannic  nitrate,  sulphate,  and  chloride  are  completely  hydrolyzed 
by  water,  while  sodium  stannate  Na^SnOa  is  comparatively  stable. 
The  dioxide  SnO2  is  an  infusible  solid,  and  resembles,  therefore, 
silicon  dioxide.  Tin  has  a  tendency  to  give  complex  acids  and 
salts,  like  H2SnCl6,  (NH4)2.SnCl6,  H2SnI6,  K2SnF6,  but  these  are 
ionized  also  to  a  small  extent  after  the  manner  of  double  salts,  giving 
ions  of  Sn '  '  '  ' .  Tin  forms  no  compounds  with  hydrogen  and  no 
salts  with  weak  acids,  like  carbonic  acid. 

Occurrence  and  Extraction.  —  Tin  has  long  been  in  use, 
specimens  of  it  being  found  in  Egyptian  tombs.  The  chief  ore  of 
tin  is  tin-stone,  or  cassiterite  SnO2,  which  consists  of  square  prismatic 
crystals  whose  dark  color  is  due  to  the  presence  of  iron  compounds. 
The  ore  is  roughly  pulverized  and  washed,  to  remove  granite  or  slate 
with  which  it  is  mixed,  and  is  then  roasted,  to  oxidize  the  sulphides 
of  iron  and  copper,  and  drive  off  the  arsenic  which  it  contains.  After 
renewed  washing  to  eliminate  sulphate  of  copper  and  oxide  of  iron, 
it  is  reduced  with  coal  in  a  reverberatory  furnace.  The  tin  is  after- 
wards remelted  at  a  gentle  heat,  and  the  pure  metal  flows  away  from 
compounds  of  iron  and  arsenic.  The  production  (1914)  was:  Straits 
Settlements  70,000  short  tons,  Bolivia  21,000,  Banka  10,400,  Cornwall 
6720,  Nigeria  (alluvial)  5600. 

Physical  and  Chemical  Properties.  —  Tin  is  a  silver-white, 
crystalline  metal  of  low  tenacity  but  great  malleability  (tinfoil).  Its 
density  is  7.3,  and  its  melting-point  about  232°.  Tin  is  dimorphous 
(p.  412).  In  1851,  the  tin  pipes  of  an  organ  were  found  to  have 
turned  largely  into  a  gray  powder.  In  1868  a  shipment  of  blocks 
of  tin  stored  in  the  custom  house  in  Petrograd  was  found  to  have 
changed  in  the  same  way.  Objects  of  tin  in  museums  frequently 
show  spots  indicating  the  presence  of  the  "tin  pest,"  as  it  was  called. 
It  now  appears  that  white,  metallic  tin  is  stable  only  above  18°,  and 


824  INORGANIC  CHEMISTRY 

that  below  this  temperature  it  is  unstable  and  is  liable  to  change 
into  gray  tin  of  lower  density  (5.85).  This  transition  point  is  similar 
to  that  of  sulphur  at  96°  (p.  412).  By  immersing  the  tin  in  a  solution 
of  pink-salt  (see  below),  the  change  is  hastened  When  the  two  kinds 
of  tin  are  used  as  the  poles  of  a  cell,  and  are  surrounded  by  pink-salt 
solution,  no  difference  in  potential  is  observed  at  -18°.  But  below 
18°,  white  tin,  being  unstable,  is  more  active  and  becomes  negative 
(giving  positive  ions),  while  above  18°,  gray  tin  becomes  negative. 

Tin-plate  (cf.  p.  800)  is  made  by  dipping  carefully  cleaned  sheets 
of  mild  steel  into  molten  tin.  Vessels  of  copper  are  also  coated, 
internally,  with  tin,  to  prevent  the  formation  of  the  basic  carbonate 
(p.  738).  For  this  purpose  they  are  cleaned  with  ammonium  chlo- 
ride, sprinkled  with  rosin  (to  reduce  the  oxide),  and  heated  to  230°. 
Molten  tin  is  then  spread  on  the  surface  with  a  piece  of  tow.  Com- 
mon pins  are  made  of  brass  wire,  and  are  coated  with  tin  by  being 
shaken  in  a  solution  containing  a  salt  of  this  metal.  The  zinc  in  the 
alloy  displaces  some  of  the  tin,  and  this  is  deposited  on  the  surface  of 
the  brass.  Alloys  of  tin,  such  as  bronze  (p.  738),  soft  solder  (50  per 
cent  lead),  pewter  (25  per  cent  lead),  and  britannia  metal  (10  per 
cent  antimony  and  some  copper),  are  much  used  in  the  arts.  On 
account  of  the  action  of  soft  water  containing  dissolved  oxygen  on 
lead  (sec  p.  830),  tin  pipes  are  preferred  for  distributing  distilled 
water  and  for  beer  pumps. 

Much  tin  is  now  recovered  by  treating  old  "tin  cans"  and  scrap 
tin-plate  with  dry  chlorine.  The  dried  gas  converts  the  tin  into 
stannic  chloride  SnCl4,  which  is  used  to  make  mordants,  but  hardly 
attacks  the  iron  (p.  222).  The  process  is  called  detinning. 

Tin,  although  it  displaces  hydrogen  from  dilute  acids,  is  not  tar- 
nished by  moist  air.  With  warm  hydrochloric  acid  it  gives  stannous 
chloride  SnCl2  and  hydrogen.  Hot,  concentrated  sulphuric  acid 
forms  stannous  sulphate  SnSO4  and  sulphur  dioxide.  Nitric  acid, 
when  cold  and  very  dilute,  interacts  with  it,  giving  stannous  nitrate, 
and  a  portion  of  the  nitric  acid  is  reduced  to  ammonia  (cf.  p.  534) : 

4Sn  +  10HNO3  ->  4Sn(NO3)2  +  3H20  +  NH4N03. 

With  concentrated  nitric  acid,  stannic  nitrate  is  formed,  but  most  of 
this  salt  is  hydrolyzed  by  the  water  at  the  high  temperature  of  the 
action  (cf.  p.  648),  and  metastannic  acid  (H2SnO3)5  (/3-stannic  acid) 
remains.  The  final  result  is  therefore  shown  by  the  equation  (sim- 
plified) : 

Sn  +  4HN03  -»  H2Sn03  +  4NO2  +  HgO. 


GERMANIUM,  TIN,  LEAD  825 

The  white,  insoluble  product  continues  to  give  nitric  acid  during  pro- 
longed washing,  and  seems  therefore  to  contain  some  basic  nitrate. 
Tin  also  displaces  hydrogen  from  caustic  alkalies,  giving  metastan- 
nates,  such  as  sodium  metastannate 


Stannous  Chloride  SnO2,2H2O.  —  This  salt  is  made  by  the 
interaction  of  tin  and  hydrochloric  acid.  When  the  colorless  crystals 
are  heated,  or  when  a  strong  aqueous  solution  is  diluted,  the  salt  is 
partially  hydrolyzed.  In  the  latter  case  the  basic  chloride  Sn(OH)Cl 
is  deposited.  By  presence  of  excess  of  hydrochloric  acid,  the  hydrol- 
ysis is  prevented.  The  solution  is  used  as  a  mordant  (p.  820). 

Stannous  chloride  tends  to  pass  into  stannic  chloride  SnCl4,  and  is 
therefore  an  active  reducing  agent.  Thus,  it  reduces  the  chlorides  of 
mercury  (p.  778)  and  of  the  noble  metals,  liberating  the  free  metals. 
The  action  is  of  the  form  Hg++  +  Sn++  ->  Hg  +  Sn++++.  Stannous 
chloride  reduces  cupric  and  ferric  chlorides  to  the  cuprous  and  ferrous 
conditions  in  like  manner: 


2FeCl3+SnCl2-*2FeCl2+SnCl4,  or  2Fe+++-f  Sn++->2Fe+++SnH 

When  stannous  chloride  solution  is  added  to  a  solution  of  sodium 
chloraurate  Na.AuCU,  the  former  reduces  the  latter,  and  metallic 
gold  is  liberated.  The  stannic  chloride  which  is  formed: 

2Na.AuCl4  +  3SnCl2  -4  2NaCl  +  2Au|  +  3SnCl4, 

is  hydrolyzed  by  the  water  (see  below),  and  the  colloidal  stannic  acid 
is  precipitated,  with  colloidal  gold  finely  dispersed  upon  it.  The 
strongly  colored  precipitate  is  called  purple  of  Cassius,  and  is  obtained 
with  different  tints,  often  red,  according  to  the  degree  of  dispersion. 
It  is  used  for  gilding  porcelain. 

Stannous  chloride  also  reduces  free  oxygen,  or,  what  is  the  same 
thing,  is  oxidized  by  the  air.  In  this  case,  stannic  chloride  is  formed 
in  an  acidified  solution  and  the  liquid  remains  clear;  in  the  neutral 
solution  a  precipitate  of  the  basic  chloride  is  formed  as  well: 

6SnCl2  +  2H20  +  O2  ->  4Sn(OH)Cl  +  2SnCl4. 

Powdered  tin,  if  placed  in  the  bottle  along  with  the  acid  solution,  will 
undo  the  effects  of  this  action  by  reducing  the  stannic  salt  to  the 
stannous  condition  once  more. 

Stannic  Chloride  SnO4.  —  When  chlorine  acts  upon  tin,  or 
upon  stannous  chloride  (either  solid  or  dissolved),  stannic  chloride  is 


826  INORGANIC  CHEMISTRY 

formed.  The  compound  is  a  colorless  liquid  (b.-p.  114°)  which  fumes 
very  strongly  in  moist  air,  giving  hydrochloric  acid  and  stannic  acid. 
It  was  formerly  known,  after  its  discoverer  (1605),  as  spiritus  fumans 
Libavii.  The  aqueous  solution,  when  freshly  made,  has  almost  no 
conductivity,  and  the  compound  is  therefore  very  slightly  ionized. 
As  hydrolysis  proceeds,  the  conductivity  increases,  but  the  hydro- 
chloric acid  is  the  conducting  substance.  After  a  time  hydrolysis 
becomes  almost  complete.  The  stannic  acid  which  is  formed  is  not 
precipitated,  however,  but  remains  in  colloidal  suspension: 

SnCl4  +  4H2O  <±  4HC1  +  Sn(OH)4. 

The  chloride,  with  small  amounts  of  water,  gives  crystalline  hydrates 
SnCl4,3H2O,  SnCl4,5H2O,  and  SnCl4,8H2O,  of  which  the  second  is 
used  as  a  mordant  under  the  name  "oxymuriate"  of  tin.  The 
double  (or  perhaps  complex)  salts  such  as  ammonium-stannic  chloride 
or  "  pink-salt"  (NH4)2SnCle  are  easily  made.  Pink  salt  is  used  as  a 
mordant  on  cotton,  and  gives  a  red  lake  with  alizarine  (p.  820). 
Stannic  bromide  SnBr4  (m.-p.  30°,  b.-p.  201°)  is  soluble  in  water. 

a-Stannic  Acid  and  its  Salts.  —  When  a  solution  of  stannic 
chloride  is  treated  with  ammonium  hydroxide,  a  white,  gelatinous  pre- 
cipitate is  formed.  To  this  the  formula  H2SnO3  is  generally  assigned  : 

SnCl4  +  4NH4OH  ->  4NH4C1  +  H2Sn03  +  H2O. 

It  is,  however,  in  reality,  amorphous,  and  loses  water  gradually  until 
the  dioxide  remains.  Thus,  neither  Sn(OH)4  nor  SnO(OH)2  is  obtain- 
able as  a  definite  compound  (p.  635).  When  stannic  oxide  is  fused 
with  caustic  soda,  a  metastannate,  namely,  the  a-stannate 
3H20,  is  formed: 

Sn02  +  2NaOH  -*  Na^nO,  +  H20. 


This  compound  is  used  as  a  mordant  under  the  name  of  "preparing 
salt."  When  its  solution  is  acidified,  a-stannic  acid,  the  actual  mor- 
dant, is  formed  by  double  decomposition.  This  a-stannic  acid  inter- 
acts readily  with  acids  and  alkalies,  and  the  chloride  obtained  from 
it  is  identical  with  stannic  chloride  described  above. 

Flannelette  and  other  cotton  goods  are  rendered  non-inflam- 
mable by  saturation  first  with  sodium  a-stannate  solution  and  then, 
after  drying,  with  ammonium  sulphate.  The  acid  is  too  feeble  and 
too  insoluble  to  form  an  ammonium  salt: 

Na2Sn03  +  (NH4)2SO4  ->  Na^SO*  +  SnO(OH)4  +  2NH3 


GERMANIUM,   TIN,   LEAD  827 

The  sodium  sulphate  is  washed  out  and  the  goods,  after  being  dried, 
contain  stannic  oxide.  The  latter  cannot  afterwards  be  removed 
by  washing,  and  the  material  is  permanently  fireproof.  Silk  is  also 
loaded  with  stannic  oxide,  the  amount  used  varying  from  25  (better 
shades  on  dyeing)  to  300  per  cent  or  more. 

The  a-stannates  of  the  metals,  aside  from  those  of  potassium  and 
sodium,  like  the  silicates  and  carbonates  which  they  much  resemble, 
are  all  insoluble  in  water,  and  may  be  made  by  double  decom- 
position. 

^-Stannic  Acid,  or  Metastannic  Acid.  —  The  product  of  the 
action  of  nitric  acid  upon  tin  is  a  hydrated  stannic  oxide  like  the  fore- 
going substance,  but  is  not  identical  with  it.  It  does  not  easily  inter- 
act with  alkalies.  By  boiling  it  with  caustic  soda,  however,  and  then 
extracting  with  pure  water,  a  soluble  sodium  /3-stannate,  Na2Sn50n, 
is  obtained.  /3-stannic  acid  is  also  very  slowly  attacked  by  acids,  and 
the  chloride  secured  from  it  is  not  identical  with  the  ordinary  chloride. 
For  these  reasons  it  is  supposed  to  be  a  hydrate  of  a  polymer  of  stan- 
nic oxide  (SnO2)5,zH2O.  When  fused  with  caustic  soda,  it  gives  the 
same  a-stannate  as  does  the  dioxide  itself. 

The  difference  between  the  properties  of  the  two  stannic  acids 
was  noticed  by  Berzelius  (1811),  and  was  the  first  case  in  which  iden- 
tity in  composition  was  found  not  to  be  accompanied  by  identity  in 
properties  (cf.  Isomers,  p.  583). 

The  Oxides  of  Tin.  —  When  stannous  oxalate  is  heated  in 
absence  of  air,  stannous  oxide  remains:  SnC2O4  — » SnO  +  CO2  +  CO. 
It  is  a  black  powder  which  burns  in  the  air,  giving  the  dioxide.  The 
corresponding  hydroxide  Sn2O(OH)2  is  formed  by  adding  sodium  car- 
bonate to  stannous  chloride  solution.  It  is  a  white  powder,  easily 
dehydrated,  and  interacts  with  alkalies  to  give  soluble  stannites, 
such  as  Na<>SnO2.  When  the  solution  is  boiled,  tin  is  deposited,  and 
sodium  stannate  is  formed,  the  behavior  resembling  that  of  cuprous 
oxide  when  heated  with  oxygen  acids  (p.  743).  With  acids,  the 
hydroxide  gives  stannous  salts. 

Stannic  oxide  SnO2  is  found  in  nature  (p.  823),  and  may  be  made 
in  pure  form  by  igniting  0-stannic  acid.  When  heated,  it  becomes 
yellow,  but  recovers  its  whiteness  when  cooled  (cf.  Zinc  oxide,  p.  770). 
Prepared  at  a  low  temperature,  it  interacts  easily  with  acids,  but  after 
strong  ignition,  is  affected  by  them  very  slowly. 


828  INORGANIC  CHEMISTRY 

The  Sulphides  of  Tin.  —  Stannous  sulphide  SnS  is  obtained 
as  a  dark-brown  precipitate  when  hydrogen  sulphide  is  led  into  a  solu- 
tion of  a  stannous  salt. 

Stannic  sulphide  SnS2  is  formed  likewise  by  precipitation,  and  is 
yellow  in  color.  It  is  made  also  by  heating  together  tin  filings, 
mercury,  sulphur,  and  ammonium  chloride.  The  mercury  and  am- 
monium chloride  are  ultimately  volatilized,  and  the  stannic  sulphide 
remains  in  the  form  of  yellow,  crystalline  scales  (" mosaic  gold"  or 
"bronze  powder").  Stannic  sulphide  loses  sulphur  when  strongly 
heated,  and  leaves  stannous  sulphide.  It  is  not  much  affected  by 
dilute  acids,  b;ut  interacts  with  solutions  of  ammonium  sulphide  (or 
sodium  sulphide),  giving  soluble  complex  sulphides,  such  as  ammo- 
nium sulphostannate : 

SnS2  +  (NH4)2S  ->  (NH4)2.SnS3. 

The  corresponding  sodium  sulphostannate  is  easily  crystallized  in  the 
form  Na2SnS3,2H2O.  Stannous  sulphide  is  not  affected  by  soluble 
sulphides,  but  polysulphides,  such  as  yellow  ammonium  sulphide, 
give  with  it  the  above  mentioned  sulphostannates : 

SnS  +  (NH4)2S  +  S  -*  (NH4)2.SnS3. 

With  acids  the  sulphostannates  undergo  double  decomposition,  but 
the  free  acid  H2.SnS3  thus  produced  is  unstable  and  breaks  up,  giving 
off  hydrogen  sulphide,  and  depositing  stannic  sulphide.* 

Analytical  Reactions  of  Salts  of  Tin.  —  The  two  ionic  forms 
of  tin,  Sn++  and  Sn4"1"1"4",  are  both  colorless.  Their  behavior  is  differ- 
ent. They  give  a  brown  and  a  yellow  sulphide,  respectively,  with 
hydrogen  sulphide.  The  interaction  of  these  sulphides  with  yellow 
ammonium  sulphide  distinguishes  them  (cf.  p.  828)  from  those  of 
cadmium,  copper,  and  other  metals  whose  sulphides  are  similarly 
inactive  towards  dilute  acids.  The  sulphides  of  arsenic,  antimony, 
and  gold  (q.v.),  however,  behave  like  those  of  tin  in  this  respect.  The 
reducing  power  of  stannous-ion  Sn4"4"  is  very  characteristic  (p.  825). 
Zinc  displaces  tin  from  solutions  of  its  salts.  The  oxides  are  reduced 
by  charcoal  in  the  reducing  part  of  the  Bunsen  flame  and  the  metal  is 
liberated. 

*  These  and  similar  compounds  are  often  called  thiostannates,  orthothian- 
timonates,  etc.  The  prefix  sulpho-  gives  more  euphonious  words,  however,  and 
is  used  here  for  all  excepting  the  thiocyanates. 


GERMANIUM,   TIN,   LEAD  &29 

LEAD  PB 

The  Chemical  Relations  of  the  Element.  —  Lead  is  both 
bivalent  and  quadrivalent.  The  oxides  PbO  and  PbO2,  and  the  cor- 
responding hydrated  oxides,  are  both  basic  and  acidic.  Lead  monox- 
ide is  a  fairly  active  base,  comparable  with  cupric  oxide,  and  lead 
dioxide  a  feeble  one.  Both  are  feebly  acidic.  The  salts  of  bivalent 
lead,  like  Pb(NO3)2,  commonly  called  the  plumbic  salts,  are  some- 
what hydrolyzed  by  water,  but  less  so  than  are  those  of  tin.  The 
tetrachloride  and  other  salts  of  quadrivalent  lead  are  completely 
hydrolyzed.  The  plumbites  Na2.PbO2  and  plumbates  Na2.PbO3  are 
hydrolyzed  to  a  considerable  extent.  All  the  compounds  in  which 
lead  is  quadrivalent  give  up  half  of  the  negative  radical  readily,  and 
are  reduced  to  the  " plumbic"  condition.  The  metal  displaces  hy- 
drogen with  difficulty,  and  is  easily  displaced  by  zinc.  Lead  com- 
pounds are  all  poisonous,  and  the  effects  of  repeated,  very  minute 
doses  are  cumulative  —  resulting  in  "lead  colic."  For  this  reason, 
the  manufacture  of  white  lead  is  forbidden  by  law  in  France,  and  is 
subject  to  strict  regulation  in  other  countries. 

Occurrence  and  Metallurgy.  —  Commercial  lead  is  almost  all 
obtained  from  galena  PbS,  which  crystallizes  in  cubes.  This  ore 
often  contains  considerable  amounts  of  silver  sulphide  Ag2S  (cf. 
p.  749),  which  is  isomorphous  with  it,  and  it  occurs  in  association 
with  sulphides  of  arsenic,  antimony,  zinc,  copper,  and  iron.  Other 
salts  of  lead  are  of  less  common  occurrence. 

The  sulphide  of  lead  is  first  roasted  until  a  sufficient  proportion 
of  it  has  been  converted  into  the  oxide  and  sulphate.  The  furnace- 
doors  are  then  closed,  and  the  temperature  raised  in  order  that  these 
products  may  interact  with  the  unchanged  part  of  the  sulphide: 

PbS  +  2PbO  ->  3Pb  +  SO2. 
PbS  +  PbSO4  ->  2Pb  +  2SO2. 

Another  plan  consists  in  heating  galenite  with  scrap  iron  or  iron  ores 
and  coal:  PbS  +  Fe  — >  Pb  +  FeS.  The  molten  ferrous  sulphide 
rises  to  the  top  as  a  matte. 

The  purification  of  the  lead  from  the  other  metals  whose  sulphides 
have  been  reduced  at  the  same  time  is  often  troublesome.  In  Parke's 
process  (p.  749)  for  the  extraction  of  the  silver  by  means  of  zinc,  the 
greater  part  of  the  foreign  metals,  with  the  exception  of  bismuth, 
pass  into  the  zinc  scum.  About  0.5  per  cent  of  zinc  remains  in  the 
lead,  and  is  oxidized  by  the  action  of  a  jet  of  steam  before  the  lead  is 


830  INORGANIC   CHEMISTRY 

poured  into  the  molds.  Lead  is  refined  electrolytically  by  the  Betts 
process.  Heavy  plates  of  the  crude  lead  form  the  anodes,  thin  sheets 
of  pure  lead  the  cathodes,  and  a  solution  of  lead  fluosilicate  PbSiF6 
the  cell  liquid.  The  operation  is  similar  to  that  for  refining  copper 
(p.  747).  Silver,  gold,  and  bismuth  are  left  as  a  sludge,  while  Zn,  Co, 
Ni,  and  Fe  go  into  solution  and  are  not  redeposited. 

The  production  (1913)  was:  United  States  412,000  short  tons, 
Spain  224,000,  Germany  199,000,  Australia  128,000. 

Physical  and  Chemical  Properties.  —  Metallic  lead  is  gray  in 
color,  very  soft,  and  of  small  tensile  strength.  Its  density  is  11.4, 
and  its  melting-point  327.4°.  While  warm,  it  is  formed  by  hydraulic 
pressure  into  pipes  which  are  used  in  plumbing  and  for  covering  elec- 
tric cables.  On  account  of  its  very  slow  interaction  with  most  sub- 
stances, sheet  lead  is  used  in  chemical  factories,  for  example,  to  line 
sulphuric-acid  chambers.  An  alloy  containing  0.5  per  cent  of  arsenic 
is  used  in  making  small  shot  and  shrapnel  bullets.  Type-metal  con- 
tains 20-25  per  cent  of  antimony  (q.v.),  and  expands  on  solidifying, 
giving  a  perfect  reproduction  of  the  mold.  In  both  cases  greater 
hardness  (cf.  p.  644)  is  secured  by  the  addition  of  the  foreign  metal. 
Solder  contains  50  per  cent  of  tin  and,  being  a  solution,  melts  at  a  low 
temperature,  and  can  be  applied  to  solid  lead  without  melting  the 
latter. 

Lead  oxidizes  very  superficially  in  the  air.  The  suboxide  Pb2O  is 
supposed  to  be  first  formed.  The  final  covering  is  a  basic  carbonate. 
Contact  with  hard  waters  confers  upon  lead  a  similar  coating  com- 
posed of  the  carbonate  and  the  sulphate.  These  deposits,  being 
insoluble  and  strongly  adherent,  enclose  the  metal  and  protect  the 
water  from  contamination  with  lead  compounds.  Pure  rain-water, 
however,  since  it  has  no  hardness,  but  contains  oxygen  in  solution, 
gives  the  hydroxide  Pb(OH)2,  which  is  noticeably  soluble.  Hence 
lead  pipes  can  safely  be  used  only  with  somewhat  hard  water.  When 
heated  in  the  air,  lead  gives  the  monoxide  PbO  or  minium  Pb3O4, 
the  latter  at  lower  temperatures. 

The  metal  displaces  hydrogen  from  hydrochloric  acid  very  slowly. 
It  is  hardly  affected  by  concentrated  sulphuric  acid  (cf.  p.  436). 
Nitric  acid  attacks  it  readily,  giving  lead  nitrate  and  oxides  of  nitro- 
gen (p.  535). 

Chlorides  and  Iodide  of  Lead.  —  Plumbic  chloride  PbCl2  is 
precipitated  when  a  soluble  chloride  is  added  to  a  solution  of  a  soluble 


GERMANIUM,   TIN,  LEAD  831 

lead  salt.  It  is  slightly  soluble  in  water  (1.5  :  100)  at  18°,  and  much 
more  so  at  100°.  In  the  saturated  solution  at  25°  about  50  per  cent 
of  the  lead  is  in  the  form  Pb++,  44  per  cent  as  PbCl+,  and  6  per  cent 
asPbCl2  (c/.  p.  439). 

Lead  tetrachloride  PbCl4  is  a  solid  at  — 15°,  and  loses  chlorine  at 
room  temperature.  It  is  made  by  passing  chlorine  into  plumbic 
chloride  suspended  in  hydrochloric  acid.  The  solution  contains 
H2PbCl6.  Ammonium  chloride  is  added  and  ammonium  chloro- 
plumbate  (NH4)2PbCl6,  analogous  to  pink-salt  (p.  826),  crystallizes 
out.  When  this  is  thrown  into  cold,  concentrated  sulphuric  acid, 
an  oil,  PbCl4,  settles  to  the  bottom.  The  oil  fumes  in  the  air  and, 
in  general,  closely  resembles  stannic  chloride  SnCl4.  With  little  water, 
it  slowly  deposits  PbCl2  and  gives  off  chlorine.  With  much  water, 
it  is  quickly  hydrolyzed,  and  lead  dioxide  is  thrown  down : 

PbCl4  +  2H20  -t  Pb02  +  4HC1. 

Lead  iodide  PbI2  (yellow)  is  formed  by  precipitation.  It  crys- 
tallizes in  yellow  scales  from  solution  in  hot  water. 

Plumbic  chloride  and  iodide  are  both  more  soluble  in  acids  or 
salts  with  a  common  negative  ion  than  they  are  in  water,  and  form 
soluble,  but  somewhat  unstable,  complex  salts. 

Oxides  and  Hydroxides.  —  There  are  five  different  oxides  of 
lead,  Pb2O,  PbO,  Pb3O4,  Pb203,  and  PbO2.  The  suboxide  Pb2O  is  a 
dark-gray  powder,  formed  by  gently  heating  the  oxalate.  Plumbic 
oxide,  or  lead  monoxide  PbO,  is  made  by  cupellation  (p.  749)  of  lead, 
and  the  solidified,  crystalline  mass  of  yellowish-red  color  is  sold  as 
litharge.  The  yellow,  powdery  form  is  called  massicot,  and  may 
be  obtained  by  heating  the  nitrate  or  carbonate.  All  the  other  oxides 
yield  this  one  when  they  are  heated  above  600°  in  the  air.  Plumbic 
oxide  takes  up  carbon  dioxide  from  the  air,  and  therefore  usually  con- 
tains a  basic  carbonate.  It  dissolves  in  warm  sodium  hydroxide  solu- 
tion, giving  a  plumbite  Na2.Pb02;  a  saturated  solution  redeposits 
part  of  the  oxide  in  crystalline  form  when  it  cools.  The  oxide  is  used 
in  making  glass  and  enamels,  and  for  preparing  salts  of  lead.  With 
glycerine,  it  gives  a  cement  for  glass  or  stone. 

Plumbic  hydroxide  Pb(OH)2  is  formed  by  precipitation.  It  gives 
up  water  in  three  stages  with  different  aqueous  tensions  (c/.  p.  654), 
the  products  in  the  order  of  decreasing  tension  being  Pb(OH)2, 
Pb20(OH)2,  Pb3O2(OH)2.  These  substances,  as  will  be  seen,  are 
equivalent  in  composition  to  PbO,H20,  2PbO,H20,  and  3PbO,H20, 


832  INORGANIC  CHEMISTRY 

respectively.     The  hydroxide  is  observably  soluble  in  water,  and 
gives  a  solution  with  a  faintly  alkaline  reaction.     With  acids  it  forms 
salts  of  lead.     It  interacts  also  with  potassium  and  sodium  hydroxides 
to  form  the  soluble  plumbites,  like  sodium  plumbite  Na2.PbO2. 
Minium,  or  red  lead,  Pb3O4,  gives  off  oxygen  when  heated: 

2Pb304  <=±  6PbO  +  02. 

The  dissociation  pressure  varies  with  the  temperature:  445°,  5  mm.; 
500°,  60  mm.;  555°,  183  mm.;  636°,  763  mm.  Since  the  partial 
pressure  of  oxygen  in  the  air  is  150  mm.,  the  substance  decomposes 
at  about  550°.  It  can  be  formed  in  air  by  reversal  of  the  action 
represented  above,  but  only  below  this  temperature,  namely  at 
470-480°  (cf.  p.  301).  In  pure  oxygen  of  one  atmosphere  pressure 
it  could  be  formed  at  600°,  but  not  at  650°.  On  account  of  unequal 
heating  during  manufacture,  commercial  red  lead  is  never  fully  oxi- 
dized, and  always  contains  litharge.  Conversely,  commercial  litharge 
usually  contains  a  little  minium. 

Minium,  when  heated  with  warm,  dilute  nitric  acid,  is  decom- 
posed, and  leaves  lead  dioxide  as  an  insoluble  powder.  Two-thirds 
of  the  lead  is  basic  and  one-third  acidic.  Minium  is  therefore  lead 
orthoplumbate  (see  below) : 

Pb2.Pb04  +  4HN03  <±  2Pb(N03)2  +  H4Pb04. 

The  double  decomposition  as  a  salt  that  it  thus  undergoes  is  followed 
by  dehydration  of  the  plumbic  acid,  which  is  unstable :  H4PbO4  — > 
Pb02  +  2H20,  and  the  dioxide  remains.  Red  lead  is  used  in  making 
flint  glass  and,  when  mixed  with  oil,  gives  a  red  paint  which  is  specially 
applicable  to  iron-work  (cf.  p.  781). 

Lead  dioxide  PbO2  may  be  obtained  as  described  above  in  the 
form  of  a  brown  powder.  Unlike  most  oxides,  it  is  a  conductor  of 
electricity.  It  is  usually  made  by  adding  bleaching  powder  to  an 
alkaline  solution  of  plumbic  hydroxide: 

Na2.Pb02  +  Ca(OCl)Cl  +  H2O  -4  2NaOH  +  CaCl2  +  Pb02|. 

In  this  action  we  may  regard  the  free  lead  hydroxide,  formed  by 
hydrolysis  of  the  plumbite,  as  being  oxidized  by  the  bleaching  powder,. 
This  dioxide  is  an  active  oxidizing  agent.  It  interacts  with,  and  sets 
fire  to,  a  stream  of  hydrogen  sulphide,  and  it  liberates  chlorine  from 
hydrochloric  acid.  With  acids  it  gives  no  hydrogen  peroxide,  and  is 
not  a  peroxidate  (p.  321).  Lead  dioxide  interacts  with  potassium  and 
,  sodium  hydroxides,  giving  soluble  plumbates.  These  are  derived 


GERMANIUM,   TIN,  LEAD  833 

from  metaplumbic  acid.  The  potassium  salt  K2PbOs,3H2O  is  analo- 
gous to  the  metastannate  K2Sn03,3H2O  (p.  826).  A  mixture  of 
calcium  carbonate  and  lead  monoxide  absorbs  oxygen  when  heated 
in  a  stream  of  air,  and  the  yellowish-red  calcium  orthoplumbate  is 
formed : 

4CaCO3  +  2PbO  +  O2  +±  2Ca2PbO4  +  4C02. 

The  action  is  reversible,  and  is  at  the  basis  of  Kassner's  plan  for 
manufacturing  oxygen  from  the  air. 

Lead  Nitrate  Pb(yVO3)2.  —  This  salt  may  be  made  by  treating 
lead,  lead  monoxide,  or  lead  carbonate  with  nitric  acid.  It  forms 
white,  anhydrous  octahedra.  The  nitrate  and  acetate  (see  below)  are 
the  salts  of  lead  which,  because  of  their  solubility,  are  most  commonly 
used.  The  solubility  of  the  nitrate  is,  48  parts  in  100  at  10°,  and 
153  parts  at  100°.  Since  the  solubility  increases  with  rise  in  tempera- 
ture, the  process  of  solution  is  accompanied  by  absorption  of  heat 
(p.  305).  On  account  of  hydrolysis,  the  solution  is  acid  in  reaction. 

Lead  Carbonate  PbCOs.  —  This  compound  is  found  in  nature 
in  rhombic  crystals,  isomorphous  with  those  of  aragonite.  It  may 
be  formed  as  a  precipitate  by  adding  a  soluble  bicarbonate  to  lead 
nitrate  solution.  With  normal  sodium  carbonate,  a  basic  carbonate 
Pb3(OH)2(CO3)2  is  deposited.  This  basic  salt  is  identical  with  white 
lead,  which,  on  account  of  its  superior  opacity,  has  better  covering 
power  than  zinc-white  (p.  770)  or  permanent  white  (p.  730).  The 
substance  is  manufactured  in  various  ways,  all  of  which  involve  the 
oxidation  of  the  lead  by  the  air,  the  formation  of  a  basic  acetate  by 
the  interaction  of  vinegar  or  acetic  acid  with  the  oxide,  and  the  sub- 
sequent decomposition  of  the  salt  by  carbon  dioxide.  The  best 
quality  is  obtained  by  the  Dutch  method.  In  this,  gratings  of  cast 
lead  (" buckles")  are  placed  above  a  shallow  layer  of  vinegar  in 
small  pots.  These  pots  are  buried  in  manure,  which  by  its  decom- 
position furnishes  the  carbon  dioxide  and  the  necessary  warmth.  The 
gratings  are  gradually  converted  into  a  white  mass  of  the  basic  car- 
bonate. The  vapor  of  acetic  acid  arising  from  the  vinegar  may  be 
regarded  as  a  catalytic  agent  (cf.  p.  436),  since  it  is  used  over  and 
over  again.  White  lead  is  made  also  by  blowing  melted  lead  into 
dust  by  means  of  steam,  beating  the  powder  with  air  and  water 
until  it  is  converted  into  the  hydrated  monoxide,  and  treating  the 
product  with  carbon  dioxide  and  vinegar. 


834  INORGANIC  CHEMISTRY 

Lead  Acetate  Pb(CO^CH^93H2O .  —  This  salt  is  made  by  the 
action  of  acetic  acid  on  litharge.  It  is  easily  soluble  in  water  and, 
from  the  sweet  taste  of  the  solution,  is  named  sugar  of  lead  (used  in 
medicine).  The  basic  salt  Pb(OH)(CO2CH3)  is  formed  by  boiling  a 
solution  of  lead  acetate  with  excess  of  litharge.  Unlike  most  basic 
salts,  this  one  is  soluble  in  water,  and  its  solution  has  a  faintly  alkaline 
reaction. 

Lead  Sulphate  PbSO*.  —  The  sulphate  occurs  in  nature  as 
anglesite,  and  is  isomorphous  with  heavy  spar.  Being  insoluble  in 
water,  it  is  easily  obtained  by  precipitation.  It  is  slightly  soluble  in 
concentrated  sulphuric  acid  (p.  436).  It  is  attacked  to  a  noticeable 
extent  by  nitric  acid,  since  this  acid  is  more  active  than  is  sulphuric 
acid  (cf.  p.  367).  It  also  interacts  with  concentrated  sodium  hydrox- 
ide solution,  on  account  of  the  removal  of  the  Pb++  ions  which  are  a 
factor  in  its  solubility  product  and  their  passage  into  the  Pb02=  anion 
of  sodium  plumbite  (cf.  p.  832).  Finally,  it  dissolves  easily  in  ammo- 
nium tartrate,  since  lead  enters  into  the  complex  anion  of  the  tartrates 
in  the  same  way  as  does  copper  (cf.  p.  744).  Barium  sulphate,  which 
is  of  the  same  order  of  insolubility  as  lead  sulphate,  is  somewhat 
affected  by  nitric  acid,  but  not  by  sodium  hydroxide  or  by  tartrates. 
The  element  barium  lacks  both  the  characteristics  which  lead  here 
exhibits. 

Lead  Sulphide  PbS.  —  Natural  lead  sulphide  (galena)  is  black, 
and  its  crystals  have  a  silvery  luster.  The  precipitated  salt  is  black 
and  amorphous.  It  is  more  easily  attacked  by  active  acids  than  is 
mercuric  sulphide  (cf.  p.  774).  Concentrated  nitric  acid,  being  an 
oxidizing  agent  as  well  as  an  acid,  interacts  with  it  readily. 

The  Storage  Battery.  —  In  the  ordinary  lead  accumulator  the 

plates  consist  of  leaden  gratings.  The  openings  are  filled  with  finely 
divided  lead  in  one  plate  and  with  lead  dioxide  in  the  other.  These, 
and  the  dilute  sulphuric  acid  in  the  cell,  are  the  active  substances 
when  the  cell  is  charged.  When  the  battery  is  used,  the  SO4=  ions 
migrate  towards  the  plates  filled  with  the  lead  (Fig.  159),  and  convert 
this  into  a  mass  of  the  insoluble  lead  sulphate :  S04~  +  Pb  — »  PbSO4 
+  2G.  These  plates  receive  the  negative  charges.  Simultaneously  ? 
the  H+  ions  move  towards  the  other  plates  and  these  reduce  to 
monoxide  the  lead  dioxide  with  which  they  are  filled : 

Pb02  +  2H+  -*  H20  +  PbO  +  20. 


GERMANIUM,  TIN,   LEAD 


835 


These  plates  acquire  positive  charges  and,  by  interaction  of  the  lead 
monoxide  with  the  sulphuric  acid,  become  filled,  like  the  negative 
plates,  with  lead  sulphate.  During  the  discharge,  much  sulphuric 
acid  is  thus  removed  from  the  cell  fluid,  and  the  approaching  exhaus- 
tion of  the  cells  can  thus  be  ascertained  by  measuring  the  specific 
gravity  of  the  fluid.  The  E.M.F.  of  the  current  is  a  little  over  2  volts. 


PbO3 


DISCHARGE 


2H+- 


PbSO4 


CHARGE 


804= 


FIG.  159. 


FIG.  160. 


The  charging  is  done  by  passing  a  current  through  the  cell,  in  the 
opposite  direction  to  the  one  which  it  yields  (Fig.  160).  The  H+  ions 
are  attracted  to  the  negative  plate  and  an  equivalent  number  of  SO4= 
ions  are  formed,  so  that  only  lead  remains  : 

PbS04  +  2H+  +  20  ->  Pb  +  2H+  +  S04= 

Simultaneously,  the  S04=  is  attracted  by  the  positive  plate  and, 
with  the  lead  sulphate  there  present,  forms  lead  persulphate  :  SO4~  -f- 
PbSO4  +  20  -»  Pb(SO4)2.  The  persulphate,  being  a  salt  of  quad- 
rivalent lead,  is  at  once  hydrolyzed  and  the  filling  of  this  plate 
is  thus  changed  into  lead  dioxide:  Pb(SO4)2  +  2H20  ->  PbO2  + 
2H2SO4.  Both  plates  are  thus  brought  back  to  the  condition  in  which 
they  were  before  the  discharge. 

The  last  set  of  changes  consumes  energy,  while  the  first  set 
liberates  energy.  Both  may  be  stated  in  a  single  equation: 


charge  —  » 

2PbSO4  +  2H20  <=»  Pb  +  2H2S04 
<—  discharge 


Pb02. 


836  INORGANIC   CHEMISTRY 

In  the  Edison  cell,  when  charged,  one  plate  is  of  iron  and  the 
other  contains  nickelic  oxide  Ni203.  The  cell  liquid  is  a  solution 
of  potassium  hydroxide.  When  the  cell  operates,  the  nickelic 
oxide  is  reduced  to  Ni(OH)2  and  the  iron  is  oxidized  to  Fe(OH)2, 
an  action  which  delivers  energy: 

Fe  +  3H2O  +  Ni2O3  <=±  Fe(OH2)  +  2Ni(OH)2. 

When  the  cell  is  charged,  the  nickel  is  reoxidized  and  the  iron  re- 
duced. 

Paints.  —  A  paint  usually  contains  three  ingredients: 

1.  The  oil  hardens  to  a  tough  resin,  being  oxidized  by  the  air 
("dries"),  and  adheres  firmly  to  the  surface  being  painted. 

2.  The  body  is  a  fine  powder  which  makes  the  paint  opaque. 
Since  the  powder  does  not  shrink,  it  also  "fills"  the  paint  and  pre- 
vents the  formation  of  minute  pores  which  otherwise  would  appear 
in  the  oil  after  drying.     White  lead  (p.  833)  is  a  common  material 
for  the  body,  but  zinc  oxide,  lithopone  (p.  730)  and  other  substances 
are  used. 

3.  Except  in  the  case  of  white  paint,  a  pigment  is  added.     Vari- 
ous oxides,  such  as  minium,  colored  salts,  and  lakes  (p.  820)  are 
used  as  coloring  matters. 

The  oil  does  not  "dry"  by  evaporation  but  gives  a  resin  by 
oxidation.  Linseed  oil  and  hemp  oil  are  commonly  used.  They 
contain  glyceryl  esters  (p.  618)  of  unsaturated  acids,  such  as  that 
of  linoleic  acid  (C3H5(CO2Ci7H3i)3),  which  contains  four  units  of 
hydrogen  less  than  stearic  acid.  The  unsaturated  part  of  the  mole- 
cule takes  up  the  oxygen.  By  previously  boiling  the  oil  with  man- 
ganese dioxide  and  other  oxides,  it  is  rendered  more  active,  and 
"dries"  more  quickly. 

Plumbers  use  a  cement  made  of  minium  and  linseed  oil,  in 
which  the  former  oxidizes  the  latter,  without  access  of  air  being 
necessary. 

Analytical  Reactions  of  Lead  Compounds.  —  Hydrogen  sul- 
phide precipitates  the  black  sulphide,  even  when  dilute  acids  are 
present.  Sulphuric  acid  throws  down  the  sulphate.  Potassium  hy- 
droxide gives  the  white  hydroxide,  which  interacts  in  excess  to  form 
the  plumbite.  Potassium  chromate  or  dichromate  (q.v.)  gives  a 
yellow  precipitate  of  lead  chromate  PbCr04,  which  is  used  as  a  pig- 
ment under  the  name  of  "chrome-yellow." 


GERMANIUM,   TIN,  LEAD  837 

TITANIUM,  ZIRCONIUM,  CERIUM,  THORIUM 

The  metals  on  the  left  side  of  the  fifth  column  of  the  periodic  table 
are  all  quadrivalent,  although  compounds  in  which  a  lower  valence 
appears  are  numerous  in  this  family.  The  first  two  are  feebly  base- 
forming  as  well  as  feebly  acid-forming;  the  last  two  are  base-forming 
exclusively. 

Titanium  occurs  in  rutile  TiTi04.  Derived  from  it  are  a  number 
of  titanates  of  the  form  K2TiO3,  titanic  iron  ore  (menaccanite)  being 
ferrous  titanate  FeTiOa. 

Zirconium  is  found  in  zircon,  the  orthosilicate  of  zirconium 
ZrSi04,  which  occurs  in  square  prismatic  crystals  isomorphous  with 
rutile,  cassiterite  (SnSnO*),  pyrolusite  (MnMn04),  and  thorite 
(ThSiO4).  The  oxide  was  used  at  one  time  in  making  the  incan- 
descent substance  in  some  forms  of  gas  lamps. 

Cerium  occurs  chiefly  in  cerite  [Ce,  La,  Nd,  Pr]Si04,H2O  (cf. 
p.  808).  The  particles  of  an  alloy  of  cerium  (70  per  cent)  and  iron 
(30  per  cent),  when  torn  off  by  a  file,  catch  fire  in  the  air.  This 
fact  is  utilized  in  making  gas-lighters  and  cigar-lighters. 

Thorium  is  found  in  thorite  ThSiO4  but  most  of  the  supply  comes 
from  monazite  sand.  The  nitrate  Th(N03)4,6H2O  is  used  in  making 
Welsbach  incandescent  mantles  (cf.  Flame,  p.  596).  The  mantle  of 
China  grass,  or  artificial  silk,  is  dipped  in  a  solution  of  this  salt 
along  with  one  per  cent  of  cerium  nitrate  Ce(NO3)4,  and  is  then 
ignited.  The  oxides  ThO2  (thoria)  and  CeO2  (ceria),  which  remain, 
form  a  fairly  coherent  mass.  Thorium  and  its  compounds  are  radio- 
active (see  Radium). 

Exercises.  —  1.  In  what  order  should  you  place  the  elements 
dealt  with  in  this  chapter,  beginning  with  the  least  metallic,  and 
ending  with  the  most  metallic  (p.  645)? 

2.  Construct  equations  showing,  (a)  the  interaction  of  tin  and 
concentrated  sulphuric  acid,  (6)  of  water  and  stannous  chloride,  (c)  of 
chlorauric  acid  and  stannous  chloride  (p.  825),  (d)  of  oxygen  and  stan- 
nous chloride  in  acid  solution,  (e)  the  decomposition  of  lead  oxalate 
(p.  831),  (/)  the  interaction  of  lead  monoxide  and  acetic  acid,  (g)  and 
of  lead  monoxide  and  lead  acetate. 

3.  To  which  class  of  ionic  actions  (pp.  402-406)  do  the  reductions 
by  stannous  chloride  belong? 

4.  What  interactions  probably  occur  when  lead  dioxide  liberates 
chlorine  from  hydrochloric  acid? 


838  INORGANIC  CHEMISTRY 

5.  How  should  you  set  about  preparing,  (a)  lead  oxalate  (insol- 
uble), (6)  lead  chlorate  (soluble)? 

6.  Should  the  formula  of  the  sulphate  of  quadrivalent  lead  be 
written  Pb(S04)2  or  PbS20s,  and  is  it  related  to  persulphuric  acid 
H2S208? 

7.  Describe  in  terms  of  the  categories  used  in  connection  with  the 
phase  rule  (p.  705)  the  system  furnished  by  minium  at  500°. 

8.  Construct  equations  for  the  formation  of  white  lead  by  the 
Dutch  process,  showing,  (1)  the  formation  of  the  basic  acetate  by 
the  action  of  oxygen,  water,  and  acetic  acid  vapor,  and  (2)  the 
action  of  carbonic  acid  on  the  product. 


CHAPTER  XLI 
ARSENIC,   ANTIMONY,   BISMUTH 

THIS  family  is  closely  related  to  the  elements  phosphorus  and 
nitrogen,  which  precede  it  in  the  same  column  of  the  periodic  table. 
In  reading  this  chapter,  therefore,  constant  reference  should  be  made 
to  the  chemistry  of  the  corresponding  compounds  of  phosphorus. 
For  a  general  comparison  of  the  elements  arsenic  (As,  at.  wt.  75), 
antimony  (Sb,  at.  wt.  120.2),  and  bismuth  (Bi,  at.  wt.  208)  with  each 
other,  and  with  the  two  already  disposed  of,  see  p.  850.  It  is  suffi- 
cient here  to  say  that  arsenic  is  mainly  an  acid-forming  element, 
and  is  therefore  non-metallic,  while  antimony  is  both  acid-forming 
and  base-forming,  and  bismuth  is  base-forming.  Each  of  the  three 
elements  gives  a  set  of  compounds  in  which  it  is  trivalent,  and 
another  in  which  it  is  quinquivalent.  None  of  the  elements,  when 
free,  displaces  hydrogen  from  dilute  acids. 

ARSENIC  As 

The  Chemical  Relations  of  the  Element.  —  Arsenic  forms  a 
compound  with  hydrogen  AsH3.  It  gives  several  halogen  derivatives 
of  the  type  AsX3  which  are  completely  hydrolyzed  by  water.  Its 
oxides  and  hydroxides  are  acidic.  Sulphates,  nitrates,  carbonates, 
and  other  salts  of  arsenic  are  not  formed.  The  complex  sulphides 
are  important.  The  soluble  compounds  of  arsenic  are  all  highly 
poisonous. 

In  many  natural  sulphides,  such  as  pyrite  FeS2  and  zinc-blende 
ZnS,  a  part  of  the  sulphur  is  replaced  by  arsenic,  which  must  here 
be  playing  the  part  of  a  bivalent  element.  When  much  arsenic  is 
present,  the  formulae  are  written:  Fe[S,AsJ2  and  Zn[S,As]. 

Occurrence  and  Preparation.  —  Arsenic  is  found  free  in 
nature.  It  occurs  also  in  combination  with  many  metals,  particularly 
in  arsenical  pyrite  (mispickel)  FeAsS.  Two  sulphides,  orpiment 
As^jSs  and  realgar  As^,  and  white  arsenic  A^Oa,  are  less  common. 

The  element  is  obtained  either  from  the  native  material  or  by 
heating  arsenical  pyrites :  FeAsS  — »  FeS  +  As.  During  the  roasting 

839 


840  INORGANIC   CHEMISTRY 

of  the  sulphur  ores  of  metals,  arsenic  trioxide  is  formed  by  the  oxida- 
tion of  the  arsenic  so  frequently  present,  and  collects  as  a  dust  in 
the  flues.  The  supply  is  greatly  in  excess  of  the  demand. 

Physical  Properties.  —  The  free  element  is  steel-gray  in  color, 
metallic  in  appearance,  and  crystalline  in  form.  When  the  vapor  is 
suddenly  cooled,  however,  a  yellow,  less  stable  variety  is  obtained, 
which  is  soluble  in  carbon  disulphide,  is  phosphorescent  in  the  air, 
and  in  other  ways  resembles  white  phosphorus. 

Elementary  arsenic  gives  off  vapor  at  180°,  and  above  600° 
acquires  a  vapor  pressure  of  760  mm.  The  density  of  the  vapor 
measured  at  644°  gives  308.4  as  the  weight  of  the  G.M.V.  (22.4  liters 
at  0°  and  760  mm.).  The  weight  combining  with  one  unit  (35.46  g.) 
of  chlorine  is  25  g.,  and  three  times  this  amount,  or  75  g.,  is  the 
smallest  weight  found  ia  the  G.M.V.  of  any  volatile  compound  of 
arsenic,  and  is  therefore  accepted  as  the  atomic  weight.  Since  308.4 
is  equal  approximately  to  4  X  75  (=  300),  the  formula  of  the  vapor 
of  the  simple  substance  at  644°  is  As4.  At  1700°  the  formula  is  As2. 

Chemical  Properties.  —  The  free  element  burns  in  the  air, 
producing  clouds  of  the  solid  trioxide  As2O3.  It  unites  directly  with 
the  halogens,  with  sulphur,  and  with  many  of  the  metals.  When 
boiled  with  nitric  acid,  chlorine  water,  and  other  powerful  oxidizing 
agents,  it  is  oxidized  in  the  same  way  as  is  phosphorus,  and  yields 
arsenic  acid  H3As(>4. 

Arsine  AsHs.  —  This  substance  corresponds  in  composition  to 
ammonia  and  phosphine,  and  some  of  the  ways  in  which  it  may  be 
formed  are  analogous  to  those  used  in  the  case  of  these  substances. 
Thus,  when  arsenic  and  zinc  are  melted  together  in  the  proportions 
to  form  zinc  arsenide  Zn3As2,  and  the  product  is  treated  with  dilute 
hydrochloric  acid,  the  result  is  similar  to  the  action  of  water  or 
dilute  acids  upon  calcium  phosphide.  Arsine  is  evolved  as  a  gas: 

Zn3As2  +  6HC1  -»  2AsH3  +  3ZnCl2. 

Arsine  (arsenuretted  hydrogen)  is  formed  also  by  the  action  of  active 
hydrogen  (cf.  p.  543)  upon  soluble  compounds  of  arsenic,  such  as 
arsenious  chloride  AsCl3  or  arsenic  acid.  When  a  solution  of  one  of 
these  substances  is  added  to  zinc  and  hydrochloric  acid  in  a  generating 
flask,  arsine  is  formed: 

AsCl3  +  3H2  -+  AsH3  +  3HC1. 


ARSENIC,   ANTIMONY,   BISMUTH  841 

This  method,  naturally,  does  not  furnish  pure  arsine,  for  free  hydrogen 
predominates  in  the  gas.  Pure  arsine  may  be  secured  by  leading  the 
mixture  with  hydrogen  through  a  U-tube  immersed  in  liquid  air.  The 
arsine  (b.-p.  —55°)  condenses  as  a  colorless  liquid  (m.-p.  —119°). 

Arsine  burns  with  a  bluish  flame,  producing  water  and  clouds  of 
arsenic  trioxide:  2AsH3  +  3O2  — >  3H2O  +  As-A.  The  combustion 
of  hydrogen  containing  arsine  produces  the  same  substances.  Since 
arsine,  when  heated,  is  readily  dissociated  into  its  constituents 
(cf.  p.  416),  the  vapor  of  free  arsenic  is  present  in  the  interior  of  the 
hydrogen  flame.  This  arsenic  may  be  condensed  in  the  form  of  a 
metallic-looking,  brownish  stain  by  interposition  of  a  cold  vessel  of 
white  porcelain.  Even  when  only  a  trace  of  the  compound  of  arsenic 
has  been  added  to  the  materials  in  the  generator,  the  stain  which  is 
produced  is  very  conspicuous.  This  behavior  thus  furnishes  us  with 
an  exceedingly  delicate  test  —  Marsh's  test  —  for  the  presence  of 
arsenic  in  any  soluble  form  of  combination.  The  compounds  of  anti- 
mony alone  show  a  similar  phenomenon  (see  Stibine).  In  carrying 
out  the  test,  a  tube  of  hard  glass  is  attached  to  the  generator,  and  is 
heated,  by  means  of  a  Bunsen  flame,  at  a  point  near  to  the  flask. 
With  this  arrangement  the  arsenic  is  deposited  in  the  form  of  a  dark, 
lustrous  ring  just  beyond  the  heated  part.  Zinc  of  special  purity 
must  be  employed  for  generating  the  hydrogen,  as  all  common  speci- 
mens of  the  metal  contain  a  sufficient  amount  of  arsenic  to  give  the 
metallic  film  without  any  actual  addition  of  an  arsenic  compound, 
and  a  blank  experiment  must  be  run,  with  other  portions  of  the  same 
reagents,  to  guard  against  the  possibility  of  its  coming  from  any  of 
them.  Arsenic  is  more  easily  detected  than  any  other  poison  which 
can  be  used  in  small  amounts. 

Arsine  is  exceedingly  poisonous,  the  breathing  of  small  amounts 
producing  fatal  effects.  It  differs  from  ammonia  more  markedly 
than  does  phosphine,  for  it  is  not  only  without  action  on  water  and 
on  acids,  but  does  not  unite  directly  even  with  the  halides  of  hydrogen. 

Halides  of  Arsenic.  —  The  halides  include  a  liquid  trifluoride 
AsF3,  a  pentafluoride,  which  is  obtained  only  as  a  double  compound 
with  potassium  fluoride,  a  liquid  trichloride  AsCl3,  a  solid  tribromide 
AsBr3,  and  a  solid  tri-iodide  AsI3. 

The  trichloride,  AsCl3,  which  is  prepared  by  passing  chlorine  gas 
into  a  vessel  containing  arsenic,  is  easily  formed  as  the  result  of  a 
vigorous  action.  It  is  a  colorless  liquid,  boiling  at  130°.  When 
mixed  with  water  it  is  at  once  converted  into  the  white,  almost  insolu- 


842  INORGANIC  CHEMISTRY 

ble  trioxide.  The  action  is  presumably  similar  to  that  of  water  upon 
the  corresponding  compound  of  phosphorus  (p.  210),  but  the  arse- 
nious  acid  for  the  most  part  loses  water  and  forms  the  insoluble  anhy- 
dride : 

AsCl3  +  3H2O  <=±  As(OH)3  +  3HC1, 
2As(OH)3<=»As203j,  +  3H2O. 

This  action,  however,  differs  markedly  from  the  other  in  that  it  is 
reversible,  and  arsenic  trioxide  interacts  with  aqueous  hydrochloric 
acid,  giving  a  solution  of  arsenious  chloride.  When  this  solution  is 
boiled,  the  volatility  of  the  arsenious  chloride  causes  it  to  be  carried 
over  with  the  hydrochloric  acid  (b.-p.  110°,  cf.  p.  563),  and  this  method 
of  separating  arsenic  from  other  substances  is  used  in  chemical 
analysis. 

Oxides  of  Arsenic.  —  Arsenic  trioxide  As2O3  is  produced  by 
burning  arsenic  in  the  air  and  during  the  roasting  of  arsenical  ores 
(p.  840),  and  is  known  as  " white  arsenic"  or  simply  " arsenic."  It  is 
purified  for  commercial  purposes  by  subliming  the  flue-dust  in  cylin- 
drical pots.  The  pure  trioxide  is  deposited  in  the  glassy  form  in 
the  upper  part  of  the  vessel.  It  passes  slowly  from  this  amorphous 
condition  into  the  common  crystalline  variety.  Its  vapor  density 
indicates  that  it  possesses  the  molecular  weight  As4O6,  but  the  simpler 
formula  expresses  its  chemical  properties  sufficiently  well. 

When  treated  with  water,  the  trioxide  dissolves  to  a  very  slight 
extent  (1.2  :  100  at  2°),  forming  arsenious  acid,  by  reversal  of  the 
second  of  the  equations  given  above.  As  usual,  the  less  stable, 
amorphous  variety  is  the  more  soluble.  In  boiling  water  the  solubil- 
ity is  much  greater  (11.5  :  100),  but  a  condition  of  equilibrium  is 
reached  very  slowly.  With  concentrated  sulphuric  acid  the  trioxide 
forms  a  sulphate  of  rather  complex  composition,  indicating  that  it 
has  basic  properties,  but  this  sulphate  is  decomposed  into  the  oxide 
and  sulphuric  acid  when  treated  with  water.  When  heated  in  a 
tube  with  carbon,  this  oxide  is  reduced,  and  the  free  element,  being- 
volatile,  is  deposited  upon  the  cold  part  of  the  tube  just  above  the 
flame.  It  is  an  active  poison,  since  it  gradually  passes  into  solution, 
forming  arsenious  acid.  The  fatal  dose  is  0.06-0.18  g.  (1-3  grains), 
but  "arsenic  eaters"  become  tolerant  of  it  and  can  take  four  times  as 
much  without  evil  effects.  It  is  contained  in  Fowler's  solution, 
which  is  used  as  a  heart  tonic. 

The  pentoxide  As205  is  a  white  crystalline  substance,  iormed  by 


ARSENIC,   ANTIMONY,  BISMUTH  843 


heating  arsenic  acid  :  2H3AsO4  —  >  AsA  +  3H20.  When  raised  to  a 
higher  temperature,  it  loses  a  part  of  its  oxygen,  leaving  the  trioxide. 
In  consequence  of  this  instability,  it  cannot  be  formed  by  direct 
union  of  oxygen  with  the  trioxide,  as  can  phosphorus  pentoxide. 

Acids  of  Arsenic.  —  When  elementary  arsenic  or  arsenious  ox- 
ide is  treated  with  concentrated  nitric  acid,  or  with  chlorine  and 
water,  orthoarsenic  acid  HsAsO4  is  produced.  The  substance  crys- 
tallizes as  a  deliquescent  white  solid  2H3AsO4,H2O.  Salts  of  this 
acid,  and  of  pyroarsenic  acid  H-iAs^y  and  metarsenic  acid  HAsO3, 
corresponding  to  the  phosphoric  acids  (p.  557),  are  known.  The 
two  last  acids,  themselves,  however,  are  not  known  as  such.  It 
has  been  shown  by  Menzies  that,  when  the  hemihydrate  of  orthoar- 
senic acid  is  dried  at  100°,  the  only  acid  obtainable  has  the  com- 
position HsAssOio  (=  5H2O,3As2O5).  When  this  acid  is  heated  more 
strongly,  it  loses  water,  leaving  the  pentoxide  A^Os.  With  the 
phosphoric  acids,  the  final  elimination  of  all  the  water  by  -simple 
heating  is  impossible.  The  chocolate-brown  silver  orthoarsenate 
Ag3AsO4  and  the  white  MgNH4As04,  like  the  corresponding  phos- 
phates, are  insoluble  in  water. 

Arsenious  acid  HsAsOs,  like  sulphurous  and  carbonic  acids,  loses 
water,  and  yields  the  anhydride,  arsenic  trioxide,  when  the  attempt  is 
made  to  obtain  it  from  the  aqueous  solution.  The  potassium  and 
sodium  arsenites,  K3AsO3  and  NasAsOs,  are  made  by  treating  arsenic 
trioxide  with  caustic  alkalies,  and  are  much  hydrolyzed  by  water. 
The  arsenites  of  the  heavy  metals  are  insoluble,  and  can  be  made  by 
precipitation.  Paris  green  and  Scheele's  green  (p.  745)  are  arsenites  of 
copper.  The  poisonous  effects  of  wall-paper  colored  with  these  com- 
pounds seem  to  be  due  to  volatile  organic  derivatives  of  arsine  which 
are  formed  by  the  action  of  a  mold.  In  cases  of  poisoning  by  white 
arsenic,  freshly  precipitated  ferric  hydroxide  or  magnesium  hydroxide 
is  administered,  since  by  interaction  with  the  arsenious  acid  they  form 
insoluble  arsenites.  The  salts  of  arsenious  acid  are  readily  oxidized, 
passing  into  arsenates.  The  action  of  a  standard  solution  of  iodine 
upon  sodium  arsenite,  for  example,  is  used  in  volumetric  analysis: 

Na3AsO3  +  H20  +  I2  ^±  Na3As04  +  2HI. 

Sulphides  of  Arsenic.  —  Three  sulphides  of  arsenic  are  known, 
As2S5,  As2S3,  As2S2.  The  first,  arsenic  pentasulphide  As^,  is  ob- 
tained as  a  yellow  powder  by  decomposition  of  the  sulpharsenates 


844  INORGANIC  CHEMISTRY 

(see  below),  and  by  leading  hydrogen  sulphide  into  a  solution  of 
arsenic  acid  in  concentrated  hydrochloric  acid.  The  latter  action 
shows  that  the  ion  As"1"1"1"1"4",  derived  from  AsCls,  is  present  in  the 
solution. 

Arsenious  sulphide  As2Ss  occurs  in  nature  as  orpiment,  and  was 
formerly  used  as  a  yellow  pigment  (auripigmentum)  .  The  word 
arsenic  is  derived  from  the  Greek  name  for  this  mineral  (d/oercvt/cov). 
It  is  obtained  as  a  citron-yellow  precipitate  when  hydrogen  sulphide 
is  led  into  an  aqueous  solution  of  arsenious  chloride. 

When  hydrogen  sulphide  is  led  into  an  aqueous  solution  of  ar- 
senious acid,  the  sulphide  is  formed,  but  remains  in  colloidal  sus- 
pension. It  is  a  negatively  charged  colloid  (p.  622),  a  small  amount 
of  H+  ion  in  the  liquid  rendering  the  whole  electrically  neutral.  It 
is  coagulated  by  adding  solutions  of  salts,  lower  concentrations 
being  sufficient  the  higher  the  valence  of  the  positive  ion  of  the  salt 
(0.05  Molar  KC1,  0.0007  M  BaCl2,  0.00009  M  A1C13). 

Realgar  As^  is  a  natural  sulphide  of  orange-red  color,  and  is  also 
manufactured  by  subliming  a  mixture  of  arsenical  pyrite  and  pyrite  : 

2FeAsS  +  2FeS2  -»  4FeS  +  As2S2T. 

It  burns  in  oxygen,  forming  arsenious  oxide  and  sulphur  dioxide,  and 
is  mixed  with  potassium  nitrate  and  sulphur  to  make  "Bengal  lights." 

Sulpharsenites  and  Sulphar  senates.  —  The  sulphides  of  ar- 
senic interact  with  solutions  of  alkali  sulphides  after  the  manner  of 
the  sulphides  of  tin  (p.  828),  giving  soluble,  complex  sulphides.  Ar- 
senious sulphide  with  colorless  ammonium  sulphide  gives  ammonium 
sulpharsenite,  and  with  the  yellow  sulphide  gives  ammonium  sulph- 
arsenate  ; 

3(NH4)2S  +  As2S3  ->  2(NH4)3AsS3, 
3(NH4)2S  +  As2S3  -f  2S  ->  2(NH4)3AsS4. 

Proustite  (p.  749)  is  a  natural  sulpharsenite  of  silver.  The  forma- 
tion of  these  soluble  compounds  is  used  in  analysis  (cf.  p.  783). 

These  salts  are  decomposed  by  acids,  and  give  free  sulpharsenious 
or  sulpharsenic  acid: 

(NH4)3.AsS3  +  3HC1  ->  3NH4C1  +  H3AsS3  ->  3H2S| 
(NH4)3.  AsS4  +  3HC1  -»  3NH4C1  +  H3AsS4  ->  3H2S  f 


These  sulpho-acids,  however,  at  once  break  up,  giving  hydrogen  sul- 
phide as  a  gas,  and  the  sulphides  of  arsenic  as  yellow  precipitates. 


ARSENIC,  ANTIMONY,   BISMUTH  845 

ANTIMONY  SB 

The  Chemical  Relations  of  the  Element.  —  Antimony  re- 
sembles arsenic  in  forming  a  hydride  SbH3  and  halides  of  the  forms 
SbX3  and  SbX5.  The  halides  are  partially  hydrolyzed  by  water  with 
ease,  but  complete  hydrolysis  is  difficult  to  accomplish  with  cold 
water.  The  oxide  Sb203  is  basic  as  well  as  feebly  acidic  (amphoteric), 
and  the  oxide  Sb2O5  is  acidic.  The  compositions  of  the  compounds 
are  similar  to  those  of  the  compounds  of  arsenic,  but  there  are  in 
addition  salts,  such  as  802(864)3,  derived  from  the  oxide  SthA.  The 
element  gives  complex  sulphides  like  those  of  arsenic. 

Occurrence  and  Preparation.  —  Antimony  occurs  free  in 
nature.  The  chief  supply,  however,  is  furnished  by  the  black  tri- 
sulphide  Sb2S3,  stibnite,  which  is  found  in  Hungary  and  Japan,  and 
forms  shining,  prismatic  crystals  of  the  rhombic  system.  Native 
stibnite  is  roasted  in  the  air  in  order  to  remove  the  sulphur,  and  the 
white  oxide  which  remains  is  mixed  with  carbon  and  reduced  by 
strong  heat: 

Sb2S3  +  5O2  74  Sb2O4  +  3SO2, 

Sb2O4+  4C 


Properties.  —  Antimony  is  a  white,  crystalline  metal,  melting 
at  630°  (b.-p.  1300°).  It  is  brittle,  and  easily  powdered.  Its  vapor 
at  1640°  has  a  density  corresponding  to  the  formula  Sb2,  while  at 
lower  temperatures  Sb4  is  present.  It  is  used  in  making  alloys  such 
as  type-metal,  stereotype-metal,  and  britannia  metal  (q.v.).  The 
alloys  of  antimony  expand  during  solidification,  and  therefore  give 
exceptionally  sharp  castings. 

Babbitt's  Metal  (Sb  3,  Zn  69,  As  4,  Pb  5,  Sn  19),  and  other  anti- 
friction alloys  used  in  lining  bearings,  contain  antimony  along  with 
zinc,  copper,  and  other  metals.  Molten  mixtures  of  metals  (al- 
loys), when  solidifying,  do  not  always  form  a  homogeneous,  solid 
mass.  In  an  anti-friction  alloy,  what  is  wanted  is  a  mass,  in  general 
soft,  but  containing  hard  particles.  The  latter  bear  most  of  the 
pressure,  yet,  as  the  alloy  wears,  they  are  pressed  into  the  softer  matrix 
so  that  a  smooth  surface  is  always  presented.  An  alloy  which  has 
the  opposite  composition,  that  is,  which  gives  a  hard  mass  con- 
taining softer  particles,  develops  heat  by  friction  much  more  rapidly. 

The  element  unites  directly  with  the  halogens.  It  does  not  rust, 
but  when  heated  it  burns  in  the  air,  forrning  the  trioxide  SbgOa  or  a 


846  INORGANIC  CHEMISTRY 

higher  oxide  Sb204.  When  heated  with  nitric  acid,  it  yields  the 
trioxide  and,  with  more  difficulty,  antimonic  acid  H3SbO4.  When 
heated  with  concentrated  sulphuric  acid,  it  forms  the  sulphate 
Sb2(S04)3. 

Stibine  SbH3.  —  The  hydride  of  antimony  SbHa  is  formed  by 
the  action  of  zinc  and  hydrochloric  acid  on  a  soluble  compound  of  anti- 
mony. By  the  action  of  dilute,  cold  hydrochloric  acid  on  an  alloy  of 
antimony  and  magnesium  (1  :  2),  a  mixture  of  hydrogen  and  stibine 
containing  as  much  as  11.5  per  cent  (by  volume)  of  the  latter  may  be 
made.  It  is  separated  by  cooling  with  liquid  air  (b.-p.  — 17°,  m.-p. 
—  88°).  It  is  more  easily  dissociated  than  is  arsine,  and  forms 'a 
deposit  of  antimony  when  a  porcelain  vessel  is  held  in  the  flame. 
The  behavior  is  in  all  respects  similar  to  that  of  arsine  (p.  841). 

The  layers  of  arsenic  or  antimony  obtained  upon  white  porcelain 
in  Marsh's  test  (p.  841)  may  be  distinguished  readily  in  several  ways. 
The  arsenic  spots  are  brownish  in  color,  lustrous,  and  volatile.  The 
antimony  spots  are  black,  smoky-looking,  and  involatile  at  the  tem- 
perature of  the  Bunsen  flame.  The  arsenic  spots  dissolve  in  a  fresh 
solution  of  bleaching  powder,  producing  calcium  chloride  and  arsenic 
acid,  while  those  of  antimony  are  unaffected.  The  arsenic  spots 
are  scarcely  attacked  by  a  solution  of  yellow  ammonium  sulphide, 
while  those  of  antimony  dissolve  readily,  forming  an  ammonium 
sulphantimoniate.  Another  distinction  between  arsine  and  stibine  is 
found  in  their  action  upon  a  solution  of  nitrate  of  silver.  Stibine  pre- 
cipitates a  silver  antimonide  Ag3Sb,  and  none  of  the  antimony  remains 
in  the  solution.  Arsine,  on  the  contrary,  precipitates  metallic  silver, 
while  arsenious  acid  remains  in  the  solution. 

Antimony  Halides.  —  The  halides  include  the  trichloride,  the 
pentachloride  SbCl5,  a  liquid  (m.-p.  -6°,  b.-p.  140°),  the  tribromide 
SbBr3,  tri-iodide  SbI3,  trifluoride  SbF3,  and  pentafluoride  SbF5. 

Antimony  trichloride  SbCl3  is  made  by  direct  union  of  the  ele- 
ments. It  forms  large,  soft  crystals,  and  used  to  be  named  "  butter 
of  antimony  (b.-p.  223°)."  When  treated  with  little  water,  it  forms 
a  white,  opaque,  insoluble  basic  salt,  antimony  oxychloride : 

SbCl3  +  H20  ^±  SbOCU  +  2HC1. 

With  a  large  amount  of  water,  a  greater  proportion  of  the  chlorine  is 
removed,  and  Sb405Cl2  (=  2SbOCl,Sb2O3)  remains.  With  boiling 
water  the  oxide  is  finally  formed.  The  action  is  not  complete  as  long 


ARSENIC,  ANTIMONY,  BISMUTH  847 

as  hydrochloric  acid  is  present.  It  may  therefore  be  reversed,  so  that 
on  addition  of  hydrochloric  acid  to  the  mixture,  a  clear  solution  of  the 
trichloride  is  re-formed.  If  the  concentration  of  the  acid  is  once 
more  reduced  by  dilution  with  water,  the  oxychloride  is  again  precip- 
itated. 

The  pentachloride  is  formed  by  leading  chlorine  over  the  tri- 
chloride. It  is  a  liquid  which  fumes  strongly  in  the  air,  being  hydro- 
lyzed  by  the  moisture. 

Oxides  of  Antimony.  —  Three  oxides  are  known:  antimony 
trioxide  Sb2O3,  antimony  pentoxide  Sb2O5,  and  an  intermediate  oxide 
Sb204.  The  trioxide  Sb2O3  is  obtained  by  oxidizing  antimony  with 
nitric  acid,  or  by  combustion  of  antimony  with  a  limited  supply  of 
oxygen.  It  is  a  white  substance,  insoluble  in  water.  It  is  in  the 
main  a  basic  oxide,  interacting  with  many  acids  to  form  salts  of  anti- 
mony. But  it  interacts  also  with  alkalies,  giving  soluble  antimonites. 
The  pentoxide  Sb20s  is  a  yellow,  amorphous  substance,  obtained  by 
heating  antimonic  acid.  It  combines  only  with  bases  to  form  salts, 
and  is  therefore  an  acid-forming  oxide  exclusively.  The  tetroxide 
Sb204  is  formed  by  heating  antimony  or  the  trioxide  in  excess  of 
oxygen.  It  is  the  most  stable  of  the  three  oxides.  It  is  neither  acid- 
nor  base-forming,  and  may  be  antimoniate  of  antimony  (SbSbO4). 

The  hydrated  trioxide  Sb(OH)3  may  be  obtained  as  a  white  pre- 
cipitate by  adding  dilute  sulphuric  acid  to  tartar-emetic  (see  below). 
It  is  insoluble,  and  easily  loses  water,  giving  the  trioxide. 

Salts  of  Antimony.  —  The  nitrate  Sb(N03)3  and  the  sulphate 

Sb2(S04)3  are  made  by  the  interaction  of  the  trioxide  with  nitric  and 
sulphuric  acids.  They  are  hydrolyzed  by  water,  giving  basic  salts, 
such  as  (SbO)2SO4  (=  Sb2O2S04),  which,  like  SbOCl,  are  derived  from 
the  hydroxide  SbO(OH) .  When  the  trioxide  is  heated  with  a  solution 
of  potassium  bitartrate  KHC4H4O6,  a  basic  salt  K(SbO)C4H406,iH2O, 
known  as  tartar-emetic,  is  formed.  This  is  a  white,  crystalline  sub- 
stance which  is  soluble  in  water  and  is  used  in  medicine.  The  univa- 
lent  group  SbO1  is  known  as  antimonyl,  and  the  above  mentioned  basic 
compounds  are  often  called  antimonyl  sulphate,  etc. 

Antimonic  Acid.  —  By  vigorous  oxidation  of  antimony  with 
nitric  acid,  or  by  decomposing  the  pentachloride  completely  with 
water,  a  white,  insoluble  substance  of  the  approximate  composition 
H3SbO4  is  obtained.  This  substance  interacts  with  caustic  potash 


848  INORGANIC   CHEMISTRY 

and  passes  into  solution.  But  the  salts  which  have  been  made  are 
pyro-  and  metantimonates.  Thus,  when  antimony  is  fused  with 
niter,  potassium  metantimonate  KSb03  is  formed.  When  dissolved, 
this  salt  takes  up  water,  and  forms  a  solution  of  the  acid  pyroanti- 
monate: 

u     2KSb03  +  H20  ->  K2H2Sb207. 

If  this  is  added  to  a  strong  solution  of  a  salt  of  sodium,  a  sodium 
pyroantimonate  Na2H2Sb207  is  thrown  down.  This  compound, 
almost  the  only  somewhat  insoluble  salt  of  sodium,  is  formed  also 
by  direct  action  of  sodium  hydroxide  upon  antimonic  acid. 

Sulphides  of  Antimony.  —  There  are  two  sulphides,  the  tri- 
sulphide  Sb2S3  and  the  pentasulphide  Sb2S5.  The  trisulphide  Sb2S3 
is  found  in  nature  as  the  black,  crystalline  stibnite.  By  the  action  of 
hydrogen  sulphide  upon  solutions  of  salts  of  antimony,  the  trisulphide 
is  precipitated  as  an  orange-red  powder,  which,  however,  after  having 
been  melted,  assumes  the  appearance  of  stibnite: 

2SbCl3  +  3H2S  <=»  Sr^Ss  J,  +  6HC1. 

The  antimony  trisulphide  is  decomposed,  and  the  above  action  is  re- 
versed, by  concentrated  hydrochloric  acid.  Like  cadmium  sulphide, 
this  substance  is  formed  only  when  the  acid  present  is  dilute. 

The  pentasulphide  Sb2S5  is  obtained  by  the  decomposition  of 
sulphantimonates  (see  below).  In  appearance  it  resembles  the  tri- 
sulphide and,  when  heated,  it  decomposes  very  readily  into  this 
substance  and  free  sulphur.  It  is  used  for  vulcanizing  rubber. 

Sulphantimonites  and  Sulphantimonates.  —  The  behavior 
of  the  sulphides  of  antimony  towards  solutions  of  the  alkali  sulphides 
is  very  similar  to  that  of  the  sulphides  of  arsenic  (p.  844).  The  tri- 
sulphide dissolves  in  colorless  ammonium  sulphide  with  difficulty, 
forming  an  unstable  ammonium  sulphantimonite : 

Sb2S3  +  3(NH4)2S  -»  2(NH4)3SbS3. 

With  the  pentasulphide,  or  with  yellow  ammonium  sulphide,  the 
action  takes  place  more  readily  and  ammonium  sulphantimonate  is 
f  omied : 

Sb2S5  +  3(NH4)2S  -4  2(NH4)3.SbS4, 
Sb2S3  +  3(NH4)2S  +  2S  ->  2(NH4)3.SbS4. 

The    most   familiar    substance    of   this    class    is    Scblippe's    salt 


ARSENIC,  ANTIMONY,   BISMUTH  849 

Na3SbS4,9H20.     Pyrargyrite  Ag3SbS3  (p.  749)  is  a  mineral  sulphan- 
timonite. 

When  acids  are  added  to  solutions  of  sulphantimonates,  the 
sulphantimonic  acid  which  is  liberated  decomposes,  and  'antimony 
pentasulphide  is  thrown  down  (see  under  Arsenic,  p.  844). 

BISMUTH  Bi 

The  Chemical  Relations  of  the  Element.  —  Bismuth  forms 
no  compound  with  hydrogen.  Its  compounds  with  the  halogens  are 
of  the  form  BiX3  and  are  hydrolyzed  by  water  giving  basic  salts. 
The  oxide  Bi2O3  is  basic  and,  although  an  oxide  Bi2Os  is  known,  it  is 
not  acidic.  Bismuth  gives  a  carbonate,  nitrate,  sulphate,  phosphate, 
and  other  salts,  in  all  of  which  it  acts  as  a  trivalent  element.  It 
forms  no  soluble  complex  sulphides. 

Occurrence  and  Physical  Properties.  —  This  element  is  found 
free  in  nature,  and  also  to  some  extent  as  trioxide  Bi2O3  and  trisul- 
phide  Bi2S3.  It  is  a  shining,  brittle  metal  with  a  reddish  tinge  (m.-p. 
271°).  Bismuth  is  one  of  the  few  substances  (see  water)  which 
expand  on  solidifying,  the  crystals  being  lighter  than  the  liquid  at 
271°.  It  is  dimorphous,  with  a  transition  point  (p.  412)  at  75°. 
When  converted  into  vapor,  its  density  at  1600-1700°  is  somewhat 
less  than  that  corresponding  to  the  formula  Bi2. 

Mixtures  of  bismuth  with  other  metals  of  low  melting-point  fuse 
at  lower  temperatures  than  do  the  separate  metals.  This  is  another 
illustration  of  the  fact  that  a  solution  melts  at  a  lower  temperature 
than  the  pure  solvent  (p.  644).  Thus,  Wood's  metal,  containing 
bismuth  (m.-p.  271°)  4  parts,  lead  (m.-p.  327°)  2  parts,  tin  (m.-p. 
232°)  1  part,  and  cadmium  (m.-p.  321°)  1  part,  melts  at  60.5°,  con- 
siderably below  the  boiling-point  of  water.  Similar  alloys  are  used 
for  safety  plugs  in  steam-boilers  and  automatic  sprinklers. 

Chemical  Properties.  —  Bismuth  does  not  tarnish,  but  when 
heated  strongly  in  the  air  it  burns  to  form  the  trioxide.  With  the 
halogens  it  forms  a  fluoride  BiF3,  a  chloride  BiCl3,  a  bromide  BiBr3, 
and  an  iodide  BiI3.  When  the  metal  is  treated  with  oxygen  acids, 
or  the  trioxide  with  any  acids,  salts  are  produced. 

Oxides.  —  In  addition  to  the  basic  trioxide,  which  is  a  yellow 
powder  obtained  by  direct  oxidation  of  the  metal  or  by  ignition  of  the 


850  INORGANIC   CHEMISTRY 

nitrate,  three  other  oxides  are  known  —  BiO,  Bi204,  and  Bi205. 
None  of  these>  however,  is  either  acid-forming,  or  base-forming. 

Salts  of  Bismuth.  —  The  salts  of  bismuth,  when  dissolved  in 
water,  like  those  of  antimony,  give  insoluble  basic  salts,  and  the 
actions  are  reversible,  the  basic  salts  being  redissolved  by  addition 
of  an  excess  of  the  acid.  In  the  case  of  the  chloride  BiCl3,H2O  and 
the  nitrate  Bi(NO3)3,5H2O,  the  actions  taking  place  are: 

BiCl3  +  2H20  +±  Bi(OH)2Cl  +  2HC1, 
Bi(N03)3  +  2H20  <=±Bi(OH)2N03  +  2HNO3. 

The  former  of  these  products,  when  dried,  loses  a  molecule  of  water, 
giving  the  oxychloride  BiOCl.  The  oxynitrate  of  bismuth  is  much 
used,  for  the  treatment  of  some  forms  of  indigestion,  under  the  name 
"subnitrate  of  bismuth."  It  is  often  contained  in  face  powders. 

It  will  be  seen  that,  although  bismuth  forms  a  colorless  ion  Bi+++, 
and  is  in  this  respect  a  metal  in  the  chemical  sense  of  the  term,  yet, 
like  many  other  metals,  it  is  related  to  the  non-metals  inasmuch  as  its 
salts  are  at  least  partially  hydrolyzed  by  water. 

The  brownish-black,  insoluble  trisulphide  Bi2S3,  may  be  obtained 
by  direct  union  of  the  elements,  or  by  precipitation  with  hydrogen 
sulphide.  On  addition  of  much  acid  the  second  of  the  above  actions 
is  reversed.  This  sulphide  is  not  affected  by  solutions  of  ammonium 
sulphide  or  of  potassium  sulphide.  It  differs,  therefore,  markedly 
from  the  sulphides  of  arsenic  and  antimony  in  its  behavior. 

THE  FAMILY  AS  A  WHOLE 

When  we  compare  the  elements  of  this  group,  taking  nitrogen  as 
the  first  of  the  family,  in  spite  of  the  fact  that  it  is  somewhat  less 
closely  related  to  the  other  members  than  they  are  to  one  another,  we 
find  an  admirable  illustration  of  the  general  principles  which  the 
periodic  system  presents. 

The  elements  themselves  change  progressively  in  physical  proper- 
ties as  the  atomic  weight  increases.  Nitrogen  is  a  gas  which  with 
sufficient  cooling  yields  a  white  solid,  phosphorus  a  white,  or  a  red 
solid,  and  arsenic,  antimony,  and  bismuth  are  metallic  in  appearance. 
The  first  combines  directly  with  hydrogen,  the  next  three  give  hy- 
drides indirectly,  and  the  last  does  not  unite  with  hydrogen  at  all. 
The  hydride  of  nitrogen  combines  with  water  to  form  a  base,  while 
the  other  hydrides  show  no  such  tendency.  Ammonia  unites  with 


ARSENIC,   ANTIMONY,   BISMUTH  851 

all  acids,  including  those  of  the  halogens,  to  form  salts;  phosphine 
with  the  hydrogen  halides  only;  the  others  do  not  combine  with 
acids  at  all.  As  regards  their  metallic  properties,  in  the  chemical 
sense,  nitrogen  and  phosphorus  do  not  by  themselves  form  positive 
ions,  and  furnish  us  therefore  with  no  salts  whatever.  Arsenic  gives 
a  trivalent  positive  ion,  which  is  found  in  solutions  of  the  halides 
only.  It  forms  no  normal  sulphates,  nitrates,  or  other  salts.  Anti- 
mony and  bismuth  both  give  trivalent  positive  ions.  The  sulphates, 
nitrates,  etc.,  of  antimony,  however,  are  readily  decomposed  by 
water  with  precipitation  of  the  hydroxide.  The  salts  of  bismuth, 
on  the  other  hand,  do  not  readily  gi^  the  pure  hydroxide  with  water, 
although  they  are  easily  hydrolyzed  to  basic  salts. 

The  halogen  compounds  of  nitrogen  and  phosphorus  are  com- 
pletely hydrolyzed  by  water,  and  do  not  exist  when  any  water  is 
present,  even  when  excess  of  the  halogen  acid  is  used.  The  halogen 
compounds  of  arsenic  are  completely  hydrolyzed  by  cold  water, 
but  exist  in  solution  in  presence  of  excess  of  the  acids.  The  halogen 
compounds  of  antimony  and  bismuth  are  incompletely  hydrolyzed 
by  cold  water. 

Each  element  gives  a  trioxide  and  a  pentoxide.  With  nitrogen 
these  are  acid-forming,  being  the  anhydrides  of  nitrous  and  nitric 
acids.  With  phosphorus  the  trioxide  and  the  pentoxide  are  anhy- 
drides of  acids.  With  arsenic  the  trioxide  is  basic  towards  the 
halogen  acids,  and  is  the  first  example  of  a  basic  oxide  which  we 
encounter  in  this  group.  The  pentoxide,  however,  is  acid-forming. 
The  trioxide  of  antimony  is  mainly  base-forming,  although  it  is 
feebly  acid-forming  also.  The  pentoxide  is  acid-forming.  The 
trioxide  of  bismuth  is  base-forming  exclusively,  and  the  pentoxide 
has  no  derivatives. 

These  statements,  which  could  easily  be  expanded,  are  sufficient 
to  show  that  when  the  periodic  law  is  borne  in  mind  it  furnishes 
valuable  aid  in  systematizing  the  chemistry  of  a  group  like  this. 

Analytical  Reactions  of  Arsenic,  Antimony ,  and  Bismuth. 

-  The  ions  which  are  most  frequently  encountered  are  As"1"4"1",  Sb+++, 
Bi+++,  AsO4~~,  and  AsO<r~.  The  first  three,  with  hydrogen  sul- 
phide, give  colored  sulphides  which  are  not  affected  by  dilute  acids. 
The  sulphides  of  arsenic  and  antimony  are  separable  from  the  sulphide 
of  bismuth  by  solution  in  yellow  ammonium  sulphide.  The  ion  of 
the  arsenates  AsO4=~  is  identified  by  its  interaction  with  salts  of 
silver  and  the  formation  of  MgNH4As04,  while  that  of  the  arsenites, 


852  INORGANIC  CHEMISTRY 

AsO3™  is  recognized  by  its  reducing  power.  Marsh's  test  (p.  841) 
enables  us  to  recognize  the  presence  of  traces  of  compounds  of  ar- 
senic and  antimony.  Oxygen  compounds  of  arsenic,  when  heated 
with  carbon,  give  a  volatile,  metallic-looking  deposit  of  arsenic. 

VANADIUM,  COLUMBIUM,  TANTALUM 

Of  these  elements,  vanadium  is  less  uncommon  than  the  others. 
It  is  found  in  rather  complex  compounds.  When  these  are  heated 
with  soda  and  sodium  nitrate,  sodium  vanadate  NaVO3  is  formed, 
and  can  be  extracted  with  water,*  Solid  ammonium  chloride  is  added 
to  the  solution,  and  ammonium  metavanadate  NH4VO3,  which  is 
less  soluble  in  solutions  of  salts  of  ammonium  (cf.  p.  698)  than  in 
water,  appears  in  the  form  of  yellow  crystals.  When  this  salt  is 
heated,  vanadic  anhydride  V2Os,  a  yellowish-red  powder,  remains. 
This  oxide  interacts  with  bases  giving  vanadates,  of  which  the  most 
stable  are  the  metavanadates.  The  element  forms  several  chlorides, 
such  as  VC12,  VC13,  VC14,  VOC13,  and  five  oxides,  V2O,  VO,  V203, 
VO2,  and  V2O5.  The  element  has  very  feeble  base-forming  properties, 
and  gives  only  a  few,  unstable  salts.  Ferrovanadium,  an  alloy  with 
iron,  is  used  in  making  vanadium  steel  (q.v.). 

Columbium  (or  niobium),  first  discovered  and  named  by  Hat- 
chett  (1801),  and  tantalum  likewise  possess  feebly  base-forming 
properties,  their  chief  compounds  being  the  columbates  and  tanta- 
lates. 

Exercises.  —  1.  How  do  you  account  for  the  fact  that  the 
molecular  weight  of  arsenic  at  644°  is  not  exactly  300,  and  why  is 
308.4  -T-  4  not  accepted  as  the  atomic  weight? 

2.  What  should  you  expect  to  be  the  interaction  of  arsine  with 
concentrated  nitric  acid? 

3.  Formulate  the  series  of  changes  involved  in  the  solution  of 
arsenic  trioxide  and  the  interaction  of  hydrochloric  acid  with  the 
arsenious  acid  so  formed  (cf.  p.  666). 

4.  What  is  the  full  significance  of  the  fact  that  arsenic  penta- 
sulphide  may  be  precipitated  by  hydrogen  sulphide  from  a  solution  of 
arsenic  acid  in  hydrochloric  acid?    Make  the  equation. 

5.  To  what  classes  of  chemical  changes  do  the  interactions  of 
arsenious  sulphide  and  antimony  trisulphide  with  yellow  ammonium 
sulphide  belong? 

6.  Construct  equations  showing  the  interaction  of,  (a)  concen- 


ARSENIC,  ANTIMONY,   BISMUTH  853 

trated  sulphuric  acid  and  antimony,  (6)  arsenic  and  bleaching-powder 
solution,  (c)  antimony  and  yellow  ammonium  sulphide,  (d)  silver 
nitrate  and  stibine,  (e)  silver  nitrate  and  arsine,  (/)  concentrated 
nitric  acid  and  antimony,  (g)  acids  and  ammonium  orthosulphanti- 
monate. 

7.   How  should  you  set  about  making  Schlippe's  salt? 


CHAPTER  XLII 
THE  CHROMIUM  FAMILY.     RADIUM 

THE  chromium  (Cr,  at.  wt.  52)  family  includes  molybdenum  (Mo, 
at.  wt.  96),  tungsten  (W,  at.  wt.  184),  and  uranium  (U,  at.  wt.  238.2), 
and  occupies  the  left  side  of  the  seventh  column  of  the  periodic  table, 
with  the  sulphur  family  on  the  right  side. 

The  Chemical  Relations  of  the  Family.  —  The  features 
which  are  common  to  the  four  elements  are  also  those  which  affiliate 
them  most  closely  with  their  neighbors  on  the  right  side  of  the  column. 
They  yield  oxides  of  the  forms  Cr03,  MoO3,  WOs,  and  UO3,  which, 
like  SO3,  are  acid  anhydrides,  and  show  the  elements  to  be  sexivalent. 
They  give  also  acids  of  the  form  H2XO4,  corresponding  to  sulphuric 
acid,  and  the  salts  resemble  the  sulphates.  Thus,  sodium  chromate 
Na2CrO4,10H2O  is  isomorphpus  with  Glauber's  salt  (p.  690),  and 
potassium  chromate  K2CrO4  with  potassium  sulphate. 

Aside  from  the  chromates,  the  first  element  forms  also  two  basic 
hydroxides  Cr(OH)2  and  Cr(OH)3,  from  which  the  numerous  chro- 
mous  (Cr++)  and  chromic  (Cr+++)  salts  are  derived.  Uranium  forms  a 
dioxide  UO2,  to  which  correspond  the  uranous  salts  like  U(SO4)2,  but 
the  mofet  familiar  salts  of  this  metal  are  basic  salts  of  the  oxide  UO3, 
and  have  the  form  UO2(NQ3)2.  Molybdenum  and  tungsten  are  not 
base-forming  elements. 

CHKOMIUM  CR 

The  Chemical  Relations  of  the  Element.  —  Chromium  gives 
four  classes  of  compounds^  and  most  of  them  are  colored  substances 
(Gk.  xpfyia,  color).  The  chromates  are  derived  from  chromic  acid 
H2CrO4,  which,  however,  is  itself  unstable,  and  leaves  the  anhydride 
CrO3  when  its  solution  is  evaporated.  The  oxide  and  hydroxide  in 
which  the  element  is  trivalent,  namely  Cr2O3  and  Cr(OH)3,  are  weakly 
basic  and  still  more  weakly  acidic.  Hence  we  have  chromic  salts  such 
as  CrCl3  and  Cr2(SC>4)3  which  are  somewhat  hydrolyzed,  but  no  carbon- 
ate, and  no  sulphide  which  is  stable  in  water.  The  compounds  in 

854 


THE   CHROMIUM   FAMILY.      RADIUM  855 

which  the  same  hydroxide  acts  as  an  acid  are  the  chromites,  and  are 
derived  from  the  less  completely  hydrated  form  of  the  oxide  CrO(OH). 
Potassium  chromite  K.CrO2  is  more  easily  hydrolyzed,  however,  than 
is  potassium  zincate  or  potassium  aluminate.  Finally,  the  chromous 
salts,  such  as  CrCl2  and  CrSO4,  correspond  to  chromous  hydroxide 
Cr(OH)2,  in  which  the  element  is  bivalent.  This  hydroxide  is  more 
distinctly  basic  than  is  chromic  hydroxide,  and  forms  a  carbonate  and 
sulphide  which  can  be  precipitated  in  aqueous  solution.  The  chro- 
mous salts  resemble  the  stannous  and  ferrous  (q.v.)  salts  in  being 
easily  oxidized  by  the  air. 

Occurrence  and  Isolation.  —  Chromium  is  found  chiefly  in 
ferrous  chromite  Fe(CrO2)2,  which  constitutes  the  mineral  chromite, 
of  which  Rhodesia  supplied  62,500  long  tons  and  New  Caledonia 
67,000  in  1915,  and  in  crocoisite  PbCKX,  which  is  chromate  of  lead. 
It  was  first  discovered  in  the  latter  mineral  by  Vauquelin  (1797). 
The  metal  is  easily  obtained  by  reduction  of  the  oxide  with  aluminium 
filings. 

Physical  and  Chemical  Properties.  —  Chromium  is  a  white, 
crystalline,  very  hard  metal  (m.-p.  1520°).  It  does  not  tarnish,  but 
when  heated  it  burns  in  oxygen,  giving  the  green  chromic  oxide  Cr203. 
It  seems  to  exist  in  two  states,  an  active  and  a  passive  one,  the  rela- 
tions of  which  are  still  somewhat  obscure.  A  fragment  which  has 
been  made  by  the  Goldschmidt  method,  or  has  been  dipped  in  nitric 
acid,  is  passive,  and  does  not  displace  hydrogen  from  hydrochloric 
acid.  When,  however,  the  specimen  is  warmed  with  this  acid,  it 
begins  to  interact,  and  thereafter  behaves  as  if  it  lay  between  zinc 
and  cadmium  in  the  electromotive  series.  If  left  in  the  air,  it  slowly 
becomes  inactive  again.  Tin  and  iron  with  hydrochloric  acid  form 
stannous  and  ferrous  chloride,  respectively,  because  the  higher 
chlorides,  if  present,  would  be  reduced  by  the  active  hydrogen.  Here, 
for  the  same  reason,  chromous  chloride  and  not  chromic  chloride  is 
formed : 

Cr  +  2HC1  -»  CrCl2  +  H2,     or     Cr  +  2H+  ->  Cr++  +  H2. 

Chromium  is  used  in  making  chrome-steel,  for  armorplate.  The 
strange  alloys,  which,  although  composed  of  active  metals,  are  not 
attacked  by  acids  (even  boiling  nitric  acid)  usually  contain  chromium 
(e.g.,  60  per  cent  Cr,  36  per  cent  Fe,  4  per  cent  Mo). 


856  INORGANIC  CHEMISTRY 

DERIVATIVES  OF  CHROMIC  ACID 

Potassium  Chromate  K2CrO4.  —  This  and  the  sodium  salt,  or 
rather  the  corresponding  dichromates  (see  below),  are  made  directly 
from  chromite,  and  form  the  starting-point  in  the  preparation  of  the 
other  compounds  of  chromium.  The  finely  powdered  mineral  is 
mixed  with  potash  and  limestone,  and  roasted.  The  lime  is  employed 
chiefly  to  keep  the  mass  porous  and  accessible  to  the  oxygen  of  the 
air,  the  potassium  compounds  being  easily  fusible: 

4Fe(Cr02)2  +  8K2CO3  +  7O2  -»  2Fe2O3  +  8K2CrO4  +  8C02. 

The  iron  is  oxidized  to  ferric  oxide,  and  the  chromium  passes  from  the 
state  of  chromic  oxide  in  the  chromite  (FeO,Cr2Os)  to  that  of  chromic 
anhydride  in  the  potassium  chromate  (K2O,CrO3).  Thus,  more 
insight  is  given  into  the  nature  of  the  action  by  the  equation  : 


4(FeO,Cr203)  +  8(K20,C02)  +  702->  2Fe203  +  8(K20,CrO3)  +  8C02. 

The  cinder  is  treated  with  hot  potassium  sulphate  solution.  This 
interacts  with  the  calcium  chromate,  which  is  formed  at  the  same 
time,  giving  insoluble  calcium  sulphate: 

CaCr04  +  K2S04  ?±  CaS04|  +  K2CrO4. 

The  whole  of  the  potassium  chromate  goes  into  solution. 

Potassium  chromate  is  pale-yellow  in  color,  rhombic  in  form  (iso- 
morphous  with  potassium  sulphate),  and  is  very  soluble  in  water 
(61  :  100  at  10°). 

Sodium  chromate  Na2CrO4,10H2O  is  made  by  using  sodium  car- 
bonate in  the  process  just  described. 

The  Dichromates.  —  When  a  solution  of  potassium  sulphate  is 
mixed  with  an  equivalent  amount  of  sulphuric  acid,  potassium  bisul- 
phate  is  obtainable  by  evaporation  :  K2S04  -f-  H2SO4  —  >  2KHS04. 
The  dry  acid  salt,  when  heated,  loses  water  (p.  437),  giving  the  pyro- 
sulphate  (or  disulphate)  :  2KHS04  +±  K2S2O7  +  H2O,  but  the  latter, 
when  redissolved,  returns  to  the  condition  of  acid  sulphate.  Now, 
when  an  acid  is  added  to  a  chromate  we  should  expect  the  chromic 
acid  H2CrO4,  thus  liberated,  to  interact,  giving  an  acid  chromate  (say, 
KHCrO4).  No  acid  chromates  are  known,  however,  and  instead  of 
them,  pyrochromates  or  dichromates  are  produced,  with  elimination 
of  water.  In  other  words,  the  second  of  the  above  actions  is  not 
appreciably  reversible  when  chromates  are  in  question: 


THE  CHROMIUM   FAMILY.      RADIUM  857 

K2Cr04  +  H2S04  -»  (H2CrO4)  +  K2S04 

K2CrO4  (+  H2CrQ4)  -» K2Cr2O7  +  H2O 

2K2Cr04  +  H2S04  -*  K2Cr207  +  H20  +  K2SO4  (1) 

In  terms  of  ions,  S2O7=  is  unstable  in  water,  and  interacts  with  the 
OH~  ion  it  contains,  giving  water  and  sulphate-ion,  while  Cr2O7=  is 
stable  in  water  and  is  formed  from  the  interaction  of  water  and 
chromate-ion : 

S207=  +  20H-  ±H?  H20  +  2S04=, 
Cr207=  +  20H-  ±=*  H20  +  2CrO4=.  (2) 

The  dichromates  of  potassium  and  sodium  are  made  by  adding 
sulphuric  acid  to  the  crude  solution  of  the  chromate  obtained  from 
chromite  (p.  856).  They  crystallize  when  the  liquid  cools,  and  the 
mother-liquor,  containing  the  potassium  sulphate  and  undeposited 
dichromate,  is  used  for  extracting  a  fresh  portion  of  cinder.  As  the 
dichromates  are  much  less  soluble  than  the  chromates,  they  crystallize 
from  less  concentrated  solutions,  and  can  therefore  be  obtained  in 
purer  condition.  Hence  the  extract  is  always  treated  for  dichromate. 

Potassium  dichromate  K2Cr207  (or  K2CrO4,CrO3)  crystallizes  in 
asymmetric  tables  of  orange-red  color.  Its  solubility  in  water  is  8: 
100  at  10°  and  12.5  :  100  at  20°.  Sodium  dichromate  Na2Cr207,2H2O 
forms  red  crystals  also,  and  its  solubility  is  109  :  100  at  15°.  This  salt 
is  now  cheaper  than  potassium  dichromate,  and  has  largely  displaced 
the  latter  for  commercial  purposes. 

By  treatment  of  the  chromates  with  larger  amounts  of  free  acid, 
other  polychromates  are  formed.  Thus,  with  increasing  amounts 
of  nitric  acid,  ammonium  chromate  gives  first  the  dichromate 
(NH4)2Cr2O7,  which  may  be  written  (NH4)2CrO4,CrO3,  then  the  tri- 
chromate  (NH4)2CrO4,2O03,  and  even  the  tetrachromate  (NH4)2- 
CrO4,3CrO3,  all  of  which  are  red  crystalline  substances. 

Chemical  Properties  of  the  Dichromates.  —  1.  When  con- 
centrated sulphuric  acid  is  added  to  a  strong  solution  of  a  dichromate 
(or  chromate),  chromic  anhydride  Cr03  separates  in  red  needles: 

Na2Cr207  +  H2S04  -*  Na^SO,  +  H20  +  2Cr03| . 

2.  Although  a  dichromate  lacks  the  hydrogen,  it  is  essentially  of 
the  nature  of  an  acid  salt,  just  as  SbOCl  lacks  hydroxyl,  but  is  essen- 
tially a  basic  salt.  Hence,  when  potassium  hydroxide  is  added  to  a 
solution  of  potassium  dichromate,  potassium  chromate  is  formed: 

K2Cr207  +  2KOH  ->  2K2Cr04  +  H20. 


858  INORGANIC  CHEMISTRY 

The  solution  changes  from  red  to  yellow,  and  the  chromate  is  obtained 
by  evaporation.  In  this  way  the  pure  alkali  chromates  are  made. 

3.  By  addition  of  potassium  dichromate  to  a  solution  of  a  salt  of  a 
metal  whose  chromate  is  insoluble,  the  chromate,  not  the  dichromate, 
is  precipitated.     This  occurs  because  there  is  always  a  little  CrQ4~ 
(equation  (2),  above)  in  the  solution  of  the  dichromate: 

2Ba(N03)2  +  K2Cr2O7  +  H2O  <=±  2BaCrO4  +  2KNO3  +  2HNO3. 

Being  essentially  an  acid  salt,  the  dichromate  produces  a  salt  and  an 
acid,  as  any  acid  salt  would  do.  For  example: 

Ba(N03)2  +  KHS04^±BaS04J.  +  KNO3  +  HNO3. 

Soluble  chromates,  naturally,  also  precipitate  insoluble  ones. 

4.  The  dichromates  of  potassium  and  sodium  melt  when  heated, 
and,  at  a  white  heat,  decompose,  giving  the  chromate,  chromic  oxide, 
and  free  oxygen.     To  make  the  equation,  we  note  that  the  dichromate, 
for  example  K2Cr2O7,  consists  of  K2CrO4  +  CrO3,  and  the  latter,  if 
alone,  will  decompose  thus :  2Cr03  — >  Cr2O3  +  3O.     Since  the  prod- 
uct must  contain  a  multiple  of  O2,  the  equation  is: 

4K2Cr207  ->  4K2CrO4  +  2Cr203  +  302. 

5.  With  free  acids,  the  dichromates  give  powerful  oxidizing  mix- 
tures, in  consequence  of  their  tendency  to  form  chromic  salts.     Since 
the  former  correspond  to  the  oxide  CrO3  and  the  latter  to  Cr2O3,  the 
passage  from  the  former  to  the  latter  must  furnish  3O  for  every 
2Cr03  transformed.     In  dilute  solutions,  unless  a  body  capable  of 
being  oxidized  is  present,  no  actual  decomposition,  beyond  the  libera- 
tion of  dichromic  acid*  occurs.     When  concentrated  hydrochloric 
acid  is  used,  this  acid  itself  suffers  oxidation: 

K2Cr2O7  +    8HC1  ->  2KC1  +  2CrCl3  +  4H2O  (+  3O) 

(3O)        +    6HC1->3H2O  +  3C12 

K2Cr207  +  14HC1  ->  2KC1  +  2CrCl3  +  7H2O  +  3C12 

When  sulphuric  acid  is  employed,  an  oxidizable  substance  such  as 
hydrogen  sulphide  (cf.  p.  418),  sulphurous  acid,  or  alcohol  must  be 
present,  if  the  dichromate  is  to  be  reduced : 

K2Cr207  +     4H2S04  r-»  K2S04  +  Cr2(S04)3  +  4H2O  (+  3O)       (1) 

(30)+      3H2S03-+3H2SO4  (2) 

or         (3O)  +  3C2H5OH  ->  3C2H40  T  +  3H20  (2') 

[alcohol]  [aldehyde] 

*  Not  shown  as  a  distinct  stage  in  the  subsequent  equations. 


THE  CHROMIUM   FAMILY.      RADIUM  859 

In  each  case  the  usual  summation  of  (1)  and  (2),  with  omission  of  the 
3O  gives  the  equation  for  the  whole  action.  When  (1)  is  dissected, 
K2O,2CrO3  giving  Cr203,3S03  +  3O  is  found  to  be  its  essential  con- 
tent. In  practice,  this  sort  of  action  is  used  for  the  purpose  of  making 
chromic  salts,  and  for  its  oxidizing  effects,  as  in  the  preparation  of 
aldehyde  and  in  the  dichromate  battery  (p.  797). 

6.  When  a  body  which  is  not  merely  oxidizable,  but  is  an  active 
reducing  agent,  is  employed,  the  dichromate  may  be  reduced  without 
the  addition  of  any  acid.  For  example,  when  warmed  with  am- 
monium sulphide,  a  dichromate  gives  chromic  hydroxide  and  free 
sulphur: 

K2Cr2O7  +  3(NH4)2S  +  7H2O->2Cr(OH)3  +  3S  +  2KOH  +  6NH4OH. 

Applications  of  Dichromates.  —  If  the  reducing  body  is  less 
active,  the  change  may  nevertheless  take  place  under'  the  influence  of 
light.  Thus,  when  paper  is  coated  with  gelatine  containing  a  soluble 
chromate  or  dichromate,  and,  after  being  dried,  is  exposed  to  light, 
chromic  oxide  is  formed  by  reduction,  and  is  adsorbed  by  the  gelatine, 
which  is  a  colloid.  This  product  will  not  swell  up  or  dissolve  in  tepid 
water,  as  does  pure  gelatine.  This  action  is  used  in  many  ways  for 
purposes  of  artistic  reproduction.  Thus,  if  the  gelatine  mixture  is 
made  up  with  lampblack,  and,  after  the  coating  has  dried,  is  covered 
with  a  negative  and  exposed  to  light,  the  parts  which  were  protected 
from  illumination  may  afterwards  be  washed  away,  while  the  carbon 
print  remains.  The  gelatine  layer  can  be  transferred  to  wood  or 
copper  before  washing.  When  materials  of  different  colors  are 
substituted  for  the  lampblack,  prints  of  any  desired  tint  may  be  made 
by  the  same  process. 

Sodium  dichromate  is  used,  instead  of  tan-bark,  in  tanning  kid 
and  glove  leathers  (chrome-tanning  process).  A  reducing  agent  is 
employed  to  precipitate  chromic  hydroxide  Cr(OH)3  in  the  leather. 
Its  use  diminishes  the  time  required  for  tanning  from  8  or  10  months 
to  a  few  hours.  The  hide  is  a  mixture  of  colloidal  materials,  and  the 
chromic  hydroxide  is  adsorbed. 

Insoluble  Chromates.  —  A  number  of  chromates,  formed  by 
precipitation  with  a  solution  of  a  soluble  chrpmate  or  dichromate,  are 
familiar.  Thus,  lead  chromate  PbCrO4  is  used  as  a  yellow  pigment. 
By  treatment  with  limewater  it  gives  a  basic  salt  of  brilliant  orange 
color  —  chrome-red  Pb2OO04.  Salts  of  calcium  give  a  yellow,  hy- 


860  INORGANIC  CHEMISTRY 

drated  calcium  chromate  CaCr04,2H20  analogous  to  gypsum  and, 
like  it,  perceptibly  soluble  in  water  (0.4  :  100  at  14°).  Barium 
chromate  BaCr04  is  also  yellow.  It  interacts  with  active  acids  to 
form  the  dichromate,  and  passes  into  solution.  Like  calcium  oxalate 
(cf.  p.  714),  it  is  not  soluble  enough  to  be  attacked  by  acetic  acid. 
Strontium  chromate  SrCrC>4,  however,  interacts  with  acetic  acid. 
Silver  chromate  Ag2CrO4  is  red,  and  interacts  easily  with  acids.  It 
will  be  observed  that  there  is  a  close  correspondence  between  the 
relative  solubilities  (see  Table)  of  the  chromates  and  the  sulphates. 

Chromyl  Chloride  CrO2O2.  —  This  compound  corresponds  to 
sulphuryl  chloride  S02C12,  and  is  made  by  distilling  a  dichromate  with 
a  chloride  and  concentrated  sulphuric  acid: 

K2Cr2O7  +  4KC1  +  3H2S04  ->  2CrO2Cl2  +  3K2SO4  +  3H20. 

The  hydrochloric  acid  liberated  from  the  chloride  may  be  supposed  to 
interact  with  chromic  acid  from  the  dichromate: 

Cr02(OH)2  +  2HC1  <=±  Cr02Cl2  +  2H20. 

Chromyl  chloride  is  a  red  liquid  (b.-p.  118°).  It  fumes  strongly  in 
moist  air,  being  hydrolyzed  by  water  (read  the  last  equation  back- 
wards). The  corresponding  bromine  and  iodine  compounds  are 
unstable,  and  when  a  bromide  or  iodide  is  treated  as  described  above, 
the  halogens  are  liberated  by  oxidation,  and  no  volatile  compound  of 
chromium  appears.  Hence,  when  an  unknown  halide  is  mixed  with 
potassium  dichromate  and  sulphuric  acid  and  distilled,  and  the  vapors 
are  caught  in  ammonium  hydroxide,  the  finding  of  a  chromate  in  the 
distillate  demonstrates  the  existence  of  a  chloride  in  the  original 
mixture: 

Cr02Cl2  +  4NH4OH  r-*  (NH4)2CrO4  +  2NH4C1  +  2H2O. 

This  action  is  used  as  a  test  for  the  presence  of  traces  of  chlorides  in 
large  amounts  of  bromides  or  iodides,  or  both. 

Chromic  Anhydride  CrO3.  —  This  oxide  is  made  as  described 
above  (par.  4,  p.  857),  and  is  often  called  " chromic  acid."  It  is 
soluble  in  water,  and  combines  with  the  latter  to  some  extent,  giving 
dichromic  acid  H2.Cr207.  In  a  solution  acidified  with  an  active  acid 
it  is  much  used  as  an  oxidizing  agent  for  organic  substances.  It 
interacts  with  acids  in  the  same  way  as  do  the  dichromates,  giving 
chromic  salts  and  furnishing  oxygen  to  the  oxidizable  body.  When 


THE   CHROMIUM   FAMILY.      RADIUM  861 

heated  by  itself,  it  loses  oxygen  readily,  leaving  the  green  chromic 
oxide:  4CrO3  ->  2Cr2O3  +  302. 

CHROMIC  AND  CHROMOUS  COMPOUNDS 

Chromic  Chloride.  —  A  hydrated  chloride  CrCl3,6H20  is  ob- 
tained by  treating  the  hydroxide  Cr(OH)3  with  hydrochloric  acid  and 
evaporating.  When  heated,  it  is  hydrolyzed,  and  chromic  oxide 
remains.  The  anhydrous  chloride  CrCl3  is  formed  by  sublimation,  as 
a  mass  of  brilliant,  reddish-violet  scales,  when  chlorine  is  led  over 
heated  metallic  chromium. 

In  this  form  the  substance  dissolves  with  extreme  slowness,  even 
in  boiling  water,  but  in  presence  of  a  trace  of  chromous  chloride,  or 
stannous  chloride  it  is  easily  soluble.  The  solution  is  green,  as  are 
all  solutions  of  chromic  salts  after  they  have  been  boiled,  but  on  stand- 
ing in  the  cold,  bluish  crystals  of  CrCl3,6H20  are  deposited.  These 
give  a  violet  solution  containing  Cr+++  -f  3C1~,  but  boiling  reproduces 
the  green  color.  The  green  material  can  also  be  obtained  in  crystals 
as  a  hexahydrate,  and  is  therefore  isomeric  (p.  583)  with  the  violet 
variety.  With  the  green  isomer,  in  cold  solution,  silver  nitrate  pre- 
cipitates at  first  only  one-third  of  the  chlorine  as, silver  chloride. 

Chromic  Hydroxide.  —  When  ammonium  hydroxide  is  added 
to  a  solution  of  a  chromic  salt,  a  hydrated  hydroxide  of  pale-blue  color, 
2Cr(OH)3,H20,  is  thrown  down.  This  loses  water  by  stages,  giving 
intermediate  hydroxides  such  as  Cr(OH)3  and  CrOOH,  and  finally 
Cr2O3.  It  interacts  with  acids,  giving  chromic  salts.  It  also  dis- 
solves in  potassium  and  sodium  hydroxides  to  form  green  solutions  of 
chromites  of  the  form  KCrO2.  When  the  solutions  of  the  alkali 
chromites  are  boiled,  the  free  hydroxide,  present  in  consequence  of 
hydrolysis,  is  converted  into  a  greenish,  less  completely  hydrated,  and 
less  soluble  variety.  This  begins  to  come  out  as  a  precipitate,  and 
soon  the  whole  action  is  reversed.  Insoluble  chromites,  such  as  that 
of  iron  Fe(Cr02)2,  are  found  in  nature.  Many  of  them,  like  Zn(CrO2)2 
and  Mg(CrO2)2,  may  be  formed  by  fusing  the  oxide  of  the  metal  with 
chromic  oxide;  the  action  being  similar  to  that  used  in  making  zinc- 
ates  (p.  771)  and  aluminates  (p.  812).  The  hydroxide  is  used  as  a 
mordant  (p.  820)  and  is  the  active  substance  in  the  chrome-tanning 
process  (p.  859). 

Chromic  Oxide  O2O3.  —  This  oxide  is  obtained  as  a  green, 
infusible  powder  by  heating  the  hydroxide,  or,  more  readily,  by  heat- 


862  INORGANIC  CHEMISTRY 

ing  dry  ammonium  dichromate;  or  by  igniting  potassium  dichromate 
with  sulphur  and  washing  the  potassium  sulphate  out  of  the  residue : 

(NH4)2Cr207  -»  N2  +  4H20  +  Cr2O3, 
K2Cr207  +  S  ->  K2S04  +  Cr2O3. 

Chromic  oxide  is  not  affected  by  acids,  but  may  be  converted  into  the 
sulphate  by  fusion  with  potassium  bisulphate.  It  is  used  for  making 
green  paint,  and  for  giving  a  green  tint  to  glass.  When  the  oxide,  or 
any  of  the  chromic  salts,  is  fused  with  a  basic  substance  such  as  an 
alkali  carbonate,  it  passes  into  the  form  of  a  chromate,  absorbing  the 
necessary  oxygen  from  the  air.  If  an  alkali  nitrate  or  chlorate  is 
added,  the  oxidation  goes  on  more  quickly.  The  alkaline  solution  of 
the  chromites  may  be  oxidized,  for  example,  by  addition  of  chlorine  or 
bromine,  and  chromates  are  formed. 

Chromic  Sulphate  O2(SO4)3,15fl"2O.  —  This  salt  crystallizes 
in  reddish-violet  crystals,  and  may  be  made  by  treating  the  hydroxide 
with  sulphuric  acid.  When  mixed  with  potassium  sulphate  solution, 
it  gives  reddish-violet,  octahedral  crystals  of  chrome-alum  (cf.  p.  813), 
K2SO4,Cr2(SO4)3,24H2O.  This  double  salt  is  most  easily  obtained  by 
reducing  potassium  dichromate  in  dilute  sulphuric  acid  by  means  of 
sulphurous  acid  (p.  858),  and  allowing  the  solution  to  crystallize. 
The  solution  of  the  crystals,  either  of  the  pure  sulphate  or  of  the  alum, 
is  bluish- violet  (Cr+++),  but  when  boiled  becomes  green.  The  green 
compound  is  formed  by  hydrolysis  and  is  gummy  and  uncrystalliz- 
able.  It  even  yields  products  which  do  not  show  the  presence  either 
of  the  Cr+++  or  the  SO4~  ion.  It  seems  to  be  formed  thus: 

2Cr2(S04)3  +  H20  <=>  Cr40(S04)4.S04  +  H2S04. 

The  green  materials  revert  slowly  to  the  violet  ones  by  reversal  of  the 
above  action  when  the  solution  remains  in  the  cold,  and  so  crystals  of 
the  sulphate  or  of  the  alum  are  obtainable  from  the  green  solutions. 

Chromic  Acetate.  —  This  salt,  Cr(CO2CH3)3,  is  made  by  treat- 
ing the  hydroxide  with  acetic  acid,  and  a  green  solution  of  it  is  used  as 
a  mordant  by  calico-printers  (cf.  p.  820). 

Chromous  Compounds.  —  By  the  interaction  of  chromium 
with  hydrochloric  acid,  or  by  reducing  chromic  chloride  in  a  stream  of 
hydrogen,  chromous  chloride  CrCl2  is  formed.  The  anhydrous  salt 
is  colorless,  and  its  solution  is  blue  (Cr++).  Like  stannous  chloride, 


THE   CHROMIUM   FAMILY.      RADIUM  863 

it  is  very  easily  oxidized  by  the  air,  a  solution  of  it  containing  excess 
of  hydrochloric  acid  being  used  in  the  laboratory  to  absorb  oxygen: 

4CrCl2  +  4HC1  +  02  ->  4CrCl3  +  2H2O. 

Chromous  hydroxide  Cr(OH)2  is  obtained  as  a  yellow  precipitate 
when  alkalies  are  added  to  the  chloride.  With  sulphuric  acid  it  gives 
chromous  sulphate  CrSO4,7H2O,  which  is  one  of  the  vitriols  (p.  771). 

Chromous  salts  give  with  ammonium  sulphide  a  black  precipitate 
of  chromous  sulphide  CrS,  and  with  sodium  acetate  a  red  precipitate 
of  chromous  acetate.  The  latter  is  not  very  soluble,  and  is  less 
quickly  oxidized  by  the  air  than  any  of  the  other  chromous  compounds. 

Analytical    Reactions    of    Chromium    Compounds.  —  The 

chromic  salts  give  the  bluish-violet  chromic-ion  Cr+++,  or  the  green 
complex  cations,  and  may  be  recognized  in  solution  by  their  color. 
The  chromates  and  dichromates  give  the  ions  CrO4=  and  Cr2O7~, 
which  are  yellow  and  red,  respectively.  From  chromic  salts,  alkalies 
and  ammonium  sulphide  precipitate  the  bluish-green  hydroxide,  and 
carbonates  give  a  basic  carbonate  which  is  almost  completely  hydro- 
lyzed  to  hydroxide.  By  fusion  with  sodium  carbonate  and  sodium 
nitrate,  they  yield  a  yellow  bead  containing  the  chromate.  The 
chromates  and  dichromates  are  recognized  by  the  insoluble  chromates 
which  they  precipitate,  and  by  their  oxidizing  power  when  mixed  with 
acids.  All  compounds  of  chromium  give  a  green  borax  bead  contain- 
ing chromic  borate,  and  this  bead  differs  from  that  given  by  com- 
pounds of  copper  (cf.  p.  746),  which  is  also  green,  in  being  unredu'cible. 

MOLYBDENUM,  TUNGSTEN,  URANIUM 

As  was  stated  at  the  opening  of  the  chapter,  these  elements  give 
acid  anhydrides  of  the  form  XOa,  and  acids  and  salts  of  the  form 
H2XO4.  They  also  give  salts  of  the  form  H2X207  corresponding  to 
the  dichromates.  Uranium  has  base-forming  properties  as  well. 

Molybdenum.  —  This  element  is  found  chiefly  in  wulfenite 
PbMoO4  and  molybdenite  MoS2.  The  latter  resembles  black  lead 
(graphite),  and  its  appearance  suggested  the  name  of  the  element 
(Gk.  poXvpSaiva,  lead).  The  molybdenite  is  converted  by  roasting 
into  molybdic  anhydride  MoO3.  When  this  is  treated  with  ammo- 
nium hydroxide,  or  with  sodium  hydroxide,  ammonium  molybdate 
(NH4)2Mo04,  or  sodium  molybdate  Na2Mo04,10H2O  is  obtained. 


864  INORGANIC  CHEMISTRY 

The  metal  itself  is  liberated  by  reducing  the  oxide  or  chloride  with  hy- 
drogen. When  pure  it  is  a  silvery  metal  and,  like  iron  (q.v.),  takes 
up  carbon  and  shows  the  phenomena  of  tempering.  The  oxides  Mo203, 
Mo02,  and  MoO3  are  known,  but  the  lower  oxides  are  not  basic. 
The  chlorides  Mo3Cl6,  MoCl3,  MoCl4,  and  MoCl6  have  been  made. 
The  chief  use  of  molybdenum  compounds  in  the  laboratory  is  in 
testing  for  and  estimating  phosphoric  acid.  When  a  little  of  a  phos- 
phate is  added  to  a  solution  of  ammonium  molybdate  in  nitric  acid, 
and  the  mixture  is  warmed,  a  copious  yellow  precipitate  of  a  phos- 
phomolybdate  of  ammonium  (NH4)3P04,llMo03,6H20  is  formed. 
The  compound  is  soluble  in  excess  of  phosphoric  acid  and  in  alkalies, 
but  not  in  dilute  mineral  acids. 

Tungsten.  —  The  minerals  scheelite  CaW04  and  wolfram 
[Fe,Mn]WO4  are  tungstates  of  calcium  and  of  iron  and  manganese, 
respectively.  By  fusion  of  wolfram  with  sodium  carbonate  and 
extraction  with  water,  sodium  tungstate  Na2WO4,2H2O  is  secured. 
It  is  used  as  a  mordant  and  for  rendering  muslin  fireproof.  Acids 
precipitate  tungstic  acid  H2WO4,H2O  from  solutions  of  this  salt.  The 
element  gives  the  oxides  W02  and  WO3,  the  latter  being  formed  by 
ignition  of  tungstic  acid.  The  chlorides  WC12,  WC14,  WC^,  and 
WCle  are  known,  the  last  being  formed  directly,  and  the  others  by 
reduction. 

The  metal  has  important  uses,  and  the  annual  production  is 
greater  than  the  total  of  all  the  metals  which  follow  it  in  the  list  on 
p.  645.  The  metal  (density  19.6)  can  be  liberated  by  reduction  of 
the  oxide  by  hydrogen  or  by  carbon.  It  has  a  higher  melting-point 
(3540°)  than  has  any  other  metal,  and,  on  this  account,  and  because 
it  is  less  volatile  than  carbon,  is  now  used  for  filaments  in  electric 
lamps.  A  carbon  filament  also  requires  3.25  watts  per  candle  power 
while  a  tungsten  filament  uses  only  1.25  watts  per  1  c.-p.  and  lasts 
twice  as  long.  The  powdered  metal  obtained  by  reduction  can  be 
pressed  into  wire  form  and  then  rolled  while  strongly  heated  by  an 
electric  current  until  a  compact  wire  is  obtained.  The  metal  can  also 
be  obtained  in  massive  form  by  reducing  the  oxide  with  aluminium, 
provided  the  crucible  and  mixture  are  heated  strongly  in  advance. 
In  1914,  in  the  United  States  alone,  about  a  hundred  million  tung- 
sten lamps  were  manufactured.  Shop  work  has  been  almost  revolu- 
tionized by  the  use  of  tungsten  steel  tools,  which  can  be  used  at 
high  speed  and,  even  when  thus  heated  red  hot  by  friction,  retain 
their  temper.  Tungsten  steel  contains  tungsten  (16  to  20  per  cent), 


THE  CHROMIUM   FAMILY.     RADIUM  865 

carbon  (0.55  to  0.75  per  cent),  chromium  (2.5  to  5  per  cent),  and 
vanadium  (0.35  to  1.5  per  cent). 

Uranium.  —  Pitchblende,  which  contains  the  oxide  UsOs  along 
with  smaller  amounts  of  many  other  elements,  is  found  mainly  in 
Joachimsthal  (Bohemia)  and  in  Cornwall.  Carnotite,  a  uranate  and 
vanadate  of  potassium  K2O,2U03,V205J3H20,  occurs  in  Colorado. 
Pitchblende  is  roasted  with  lime,  the  calcium  uranate  CaUO4  thus 
formed  is  decomposed  with  sulphuric  acid,  giving  uranyl  sulphate 
UO2SO4.  When  excess  of  sodium  carbonate  is  added  to  the  solution 
of  the  latter,  the  foreign  metals  are  precipitated  and  sodium  diuranate 
Na2U2O7,7H20,  which  is  also  thrown  down,  dissolves  in  the  excess  as 


After  nitration,  the  diuranate  of  sodium  is  reprecipitated  by 
neutralizing  with  sulphuric  acid  and  boiling.  This  salt  is  used  in 
making  uranium  glass,  which  shows  a  yellowish-green  fluorescence. 
The  property  is  due  to  the  fact  that  the  wave-length  of  part  of  the 
invisible,  ultra-violet  rays  of  the  sunlight  are  lengthened,  and  a 
greenish  light  is  therefore  in  excess.  The  oxides  are  UO2  a  basic  oxide, 
U20s,  UsOs  the  most  stable  oxide,  UOs  uranic  anhydride,  and  UO4  a 
peroxide. 

When  the  oxide  U02  is  treated  with  acids,  it  gives  uranous  salts 
such  as  uranous  sulphate  U(SO4)2,4H2O.  Uranic  anhydride  and 
uranic  acid  interact  with  acids,  giving  basic  salts,  such  as  UO2SO4,- 
3|H2O,  and  UO2(NO3)2,6H2O,  which  are  named  uranyl  sulphate, 
uranyl  nitrate,  and  so  forth.  They  are  yellow  in  color,  with  green 
fluorescence.  Ammonium  sulphide  throws  down  the  brown,  unstable 
uranyl  sulphide  UO2S  from  their  solutions. 

RADIOACTIVE  ELEMENTS 

Historical.  —  We  have  seen  (p.  469)  that  in  an  evacuated  tube, 
through  which  an  electric  discharge  is  passed,  the  "rays"  emanat- 
ing from  the  cathode  (cathode  rays)  strike  the  anti-cathode  and 
the  glass  beyond  it.  Cathode  rays  were  discovered  by  Sir  William 
Crooks  (1878),  and  later  were  shown  to  consist  of  particles  of  nega- 
tive electricity  or  electrons,  each  having  a  mass  about  y*W  of  that 
of  an  atom  of  hydrogen.  They  produce  in  the  glass  a  greenish- 
yellow,  fluorescent  light.  Rontgen  (1895)  accidentally  discovered 
that  this  light  (X-rays)  penetrated  paper,  flesh,  and  other  materials 
composed  of  elements  of  low  atomic  weight  and  acted  upon  photo- 


866  INORGANIC  CHEMISTRY 

graphic  plates.  In  1896  Henri  Becquerel  observed  that  minerals 
containing  uranium  gave  off  a  sort  of  radiation  which  could  pene- 
trate black  paper,  that  was  opaque  to  ordinary  light,  and  reduce 
the  silver  bromide  on  a  photographic  plate  placed  beneath  the 

papers.  He  also  discovered  that  an  elec- 
trometer (Fig.  161),  in  which  the  gold  leaves 
had  been  caused  to  separate  by  charging 
with  electricity,  lost  its  charge  rapidly  when 
the  uranium  ore  (or  salt)  was  brought  near 
(3-4  cm.)  to  the  knob  connected  with  the 
leaves.  The  uranium  material  rendered  the 
air  a  conductor  (" ionized"  the  air)  and 
this  effect  permitted  the  escape  of  the  elec- 
FlQ  161  trie  charge,  which  otherwise  would  have 

been  retained  for  a  considerable  time.     In 

the  quantitative  measurement  of  radioactivity,  we  now  compare  the 
times  required  for  the  discharge  of  an  electroscope  by  different 
specimens  of  radioactive  matter.  The  presence  of  10~12  g.  of  such 
matter  can  thus  be  detected. 

The  radioactivity  of  every  pure  uranium  compound  is  propor- 
tional to  its  uranium  content.  The  ores  are,  however,  relatively 
four  times  as  active.  This  fact  led  M.  and  Mme.  Curie,  just  after 
1896,  to  the  discovery  that  the  pitchblende  residues,  from  which 
practically  all  of  the  uranium  had  been  extracted,  were  neverthe- 
less quite  active.  About  a  ton  of  the  very  complex  residues  having 
been  separated  laboriously  into  the  components,  it  was  found  that 
a  large  part  of  the  radioactivity  remained  with  the  sulphate  of  barium. 
From  this  barium  sulphate,  a  product  free  from  barium,  and  at  least 
one  million  times  more  active  than  uranium,  was  finally  secured  in 
the  form  of  the  bromide.  The  nature  of  the  spectrum  and  the  chemi- 
cal relations  of  the  element,  now  named  radium,  placed  it  with  the 
metals  of  the  alkaline  earths.  The  ratio  by  weight  of  chlorine  to 
radium  in  the  chloride  is  35.46  :  113,  so  that,  on  the  assumption  that 
the  element  is  bivalent,  its  chloride  is  RaCl2  and  its  atomic  weight  is 
226.  With  this  value  it  occupies  a  place  formerly  vacant  in  the 
periodic  table. 

In  1910  Mme.  Curie  obtained  metallic  radium  by  electrolyzing 
a  solution  of  radium  chloride,  using  a  mercury  cathode,  and  ex- 
pelling the  mercury  by  distillation.  It  was  a  white  metal  (m.-p. 
700°)  which,  like  calcium,  quickly  tarnished  in  the  air  and  displaced 
hydrogen  from  water. 


THE   CHROMIUM   FAMILY.      RADIUM  867 

The  Nature  of  the  "Rays."  —  Many  properties  show  that 
the  "rays"  emitted  by  compounds  of  uranium  and  of  radium  are 
of  three  kinds.  They  are  most  sharply  distinguished  from  one 
another  when  allowed  to  pass  through  a  powerful  magnetic  field. 
The  alpha-rays  are  positively  charged  and  are  bent  in  one  direction 
while  the  beta-rays  are  negative  and  are  bent  in  the  other.  The 
gamma-rays  are  not  affected. 

The  alpha-rays  are  atoms  of  helium  (p.  511)  thrown  off  in  straight 
lines  with  varying  initial  velocities,  averaging  about  one-tenth  that 
of  light  (say,  30,000  kilometers  per  second.  The  a-particles  from 
Ra-C,  e.g.,  19,220  kilom.  per  sec.).  Each  such  atom  bears  a  double 
positive  charge  (the  unit  being  the  charge  on  a  univalent  positive 
ion),  and  a  delicate  electroscope  readily  indicates  the  entrance  of  a 
single  atom.  These  alpha-particles,  being  each  four  times  as  heavy 
as  an  atom  of  hydrogen,  plough  their  way  through  tens  of  thousands 
of  air-molecules  and  usually  go  about  3-8  cm.  be- 
fore being  stopped.  The  emission  of  atoms  of 
helium  can  be  detected  by  means  of  Crookes  spin- 
thariscope (Fig.  162).  The  particle  of  radium 


bromide  is  at  B,  and  some  of  the  charged  helium  FlG  162 

atoms  strike  a  surface  C  covered  with  zinc  sul- 
phide, producing  faint  flashes  of  light.  The  lens  A  magnifies  the 
flashes  and  the  latter  can  be  seen  in  a  dark  room  after  the  eye  has 
become  thoroughly  rested  (15-20  minutes).  The  helium  gas  given 
off  by  radium  compounds  was  collected  by  Soddy,  working  in  Ram- 
say's laboratory  and  identified,  and  its  rate  of  production  was  meas- 
ured. The  amount  was  equal  to  158  cubic  mm.  per  1  g.  of  radium 
per  year. 

The  alpha-particles,  in  passing  through  the  air-molecules,  ionize 
the  air,  and  the  ionized  air  has  the  same  power  that  dust  possesses 
(p.  505)  of  affording  nuclei  on  which  moisture  can  condense.  Hence, 
when  a  particle  of  a  radium  compound  is  supported  in  a  flask  con- 
taining air  saturated  with  moisture,  and  the  air  is  suddenly  cooled  by 
expansion,  the  paths  of  the  particles  become  lines  of  fog.  With 
powerful  illumination,  the  fog-tracks  (Fig.  163)  can  be  photographed 
(C.  T.  R.  Wilson),  and  the  lengths  of  the  paths  can  be  measured. 

The  beta-particles  are  electrons  (p.  354),  or  unit  charges  of  nega- 
tive electricity,  and  are  shot  out  with  a  velocity  approaching  that 
of  light  (300,000  kiloms.  per  sec.).  They  are  therefore  identical 
with  cathode  rays,  but  move  many  times  more  rapidly.  Being  very 
light  (weight,  y-gViF  of  an  atom  of  hydrogen),  their  paths,  although 


868 


INORGANIC  CHEMISTRY 


FIG.  163. 


THE  CHROMIUM   FAMILY.      RADIUM  869 

straight  at  first,  soon  become  tortuous  owing  to  collisions  with  the 
relatively  massive  air-molecules.  Half  of  them  are  lost  after  going 
about  4  cm.  Their  fog  tracks  are  fainter  than  are  those  of  the  a- 
particles  and  extremely  tangled.  Being  much  lighter  than  a-particles, 
their  paths  are  actually  coiled  into  circles  or  spirals  by  a  magnetic  field. 

The  gamma-rays  are  identical  with  X-rays  (vibrations  in  the 
ether  of  short  wave-length,  p.  469),  and  are  produced,  like  the  latter, 
by  the  impacts  of  the  electrons  on  the  surrounding  matter. 

The  helium  atoms  are  almost  all  stopped  by  a  sheet  of  paper  or 
by  aluminium  foil  0.1  mm.  thick.  The  electrons  have  greater 
penetrating  power,  many  passing  through  gold-leaf,  but  being  prac- 
tically all  stopped  by  a  sheet  of  aluminium  1  cm.  thick.  The  gamma- 
rays  (X-rays),  however,  are  able  to  penetrate  relatively  thick  layers 
of  metals  and  other  materials  of  low  atomic  weight. 

One  of  the  most  striking  facts  is  that  the  stoppage  by  the  air  of 
so  many  rapidly  moving  particles  results  in  the  production  of  much 
heat.  One  gram  of  radium  would  produce  about  120  cal.  per  hour. 

Disintegration.  —  The  emission  of  atoms  of  helium  and  of 
electrons  was  first  explained  by  Rutherford  (1902-3),  then  of  McGill 
University,  Montreal,  as  being  due  to  the  spontaneous  disintegration 
of  the  atoms  of  uranium,  radium,  and  other  radioactive  elements. 
Thus,  Rutherford  was  the  first  to  show  that  radium  compounds 
produced  a  gaseous  substance  called  the  radium  emanation  (niton), 
which  was  the  residue  left  after  the  emission  of  one  atom  of  helium 
from  an  atom  of  radium.  This  gas  was  itself  radioactive  and  under- 
went further  disintegration,  depositing  a  solid  radioactive  residue 
on  bodies  in  contact  with  it.  Furthermore,  every  known  uranium 
ore  contains  radium  (McCoy)  and  radium  emanation  (Boltwood)  in 
amounts  proportional  to  the  uranium  content.  Also,  after  the  radium 
has  been  remored,  the  pure  uranium  compound  gives  off  at  first  only 
a-particles,  but  gradually  recovers  its  whole  radioactivity  and  is  then 
found  to  contain  radium  emanation  once  more  (Soddy).  It  thus 
appears  that  uranium  is  the  starting  point,  and  that  the  disintegra- 
tion proceeds  by  steps,  producing  a  number  of  different  products. 
Each  of  these  is  formed  from  one  such  product  and  by  disintegration 
furnishes  another. 

Unlike  ordinary  chemical  change,  the  rate  of  disintegration  is 
not  affected  by  conditions.  It  can  neither  be  started  nor  stopped  at 
will.  It  is  no  more  vigorous  at  2000°  than  at  -200°.  Other 
changes  occur  between  atoms,  these  within  each  atom. 


870  INORGANIC   CHEMISTRY 

The  law,  due  also  to  Rutherford,  describing  the  rate  at  which 
any  one  radioactive  element  disintegrates  is  simple.  Only  a  certain 
fraction  of  the  whole  of  any  one  specimen  undergoes  the  change  in 
unit  time.  Thus,  as  the  total  amount  diminishes  because  of  the 
change,  the  amount  changing  during  the  next  unit  of  time,  being 
a  constant  fraction  of  the  whole,  must  be  less.  Hence  an  infinite 
time  would  be  required  for  the  complete  disintegration  of  any  one 
specimen.  For  convenience,  therefore,  it  is  sometimes  the  custom 
to  give  as  a  specific  property  of  each  radioactive  element  the  time 
required  for  the  decay  of  half  its  amount  and  therefore  the  loss  of 
half  of  its  radioactivity.  More  usually,  the  property  given  is  the 
one  called  the  average  life  of  the  element.  The  value  of  this  is  equal 
to  the  inverse  of  the  fraction  disintegrating  per  unit  time,  and  is 
about  1.44  times  the  period  of  half  change.  Numerically  it  is  the 
sum  of  the  separate  periods  of  future  existence  of  all  the  atoms 
divided  by  the  number  of  such  atoms  present  at  the  starting  point. 

Radium  emits  helium  atoms  at  the  rate  of  3.4  X  1010  per  gram 
per  second.  From  this  fact,  we  can  calculate  its  average  life  to 
be  about  2400  years.  Hence,  if  it  were  not  continuously  being 
produced  (from  uranium),  the  whole  supply  would  have  been  ex- 
hausted long  before  the  earth  reached  a  habitable  condition. 

The  Uranium  Group  of  Radioactive  Elements.  —  The  fol- 
lowing shows  the  various  elements  produced  from  uranium  by  suc- 
cessive disintegrations.  When  a  helium  atom  or  an  electron  is 
expelled,  the  fact  is  shown  by  the  symbols  He  and  e,  respectively. 
The  first  number  below  each  element  is  the  average  life  of  that 
member  of  the  series  (y.  =  year,  d.  =  day,  h.  =  hour,  m.  =  minute, 
s.  =  second).  The  second  number  is  the  atomic  weight,  obtained 
by  subtracting  from  the  at.  wt.  of  uranium  (238.2)  the  weight  (4) 
of  each  helium  atom  emitted. 

Ui       -»He  +  U-Xi      -»e  +  U-X2-»€   +  U2          -» He  +  Ionium 

8X10»y.  35.5  d.  1.65m.  3X10"  y.  2X10*  y. 

238.2  234.2  234.2  234.2  230.2 

->He  +  Ra     ->    He  +  Niton  -»He  +  Ra-A   -»He  +  Ra-B 

2440  y.  5.55  d.  4.3m.  38.5m. 

226  222  218  214 

-»e     +Ra-C->    c     +Ra-Ci  ->He  +  Ra-D   ->  e     +  Ra-E 

28.1m.  10-«s.  24  y.  7.2  d. 

214  214  210  210 

-»€    +Ra-F-»    He  +  Pb  (end) 

196  d. 

210  206 


THE   CHROMIUM   FAMILY.      RADIUM  871 

A  purified  salt  of  uranium  recovers  half  its  activity  in  about 
three  weeks,  and  reaches  full  equilibrium  in  from  six  months  to 
a  year.  An  equilibrium  is  attained  when  the  speed  at  which  each 
disintegration  product  is  being  formed  is  balanced  by  the  equal 
speed  with  which  it  is  passing  into  the  next  member  of  the  series. 
The  complex  operations  required  for  studying  all  the  members 
of  the  series  cannot  be  given  here.  It  may  be  said,  however,  that 
a  pure  uranium  salt  in  solution  gives  with  ammonium  carbonate 
a  precipitate  which  is  wholly  soluble  in  excess  of  the  reagent.  After 
about  a  year,  another  portion  of  the  same  specimen  leaves  a  slight 
precipitate  which  is  insoluble  in  excess  and  contains  the  products  of 
disintegration,  chiefly  U-X,  which  was  first  obtained  in  this  way  by 
Crookes. 

The  radium  emanation  was  shown  by  Ramsay  to  be  one  of  the 
inert  gases  (p.  511),  and  was  renamed  niton.  Its  density  was  de- 
termined experimentally  with  a  small  sample,  using  a  micro-balance 
capable  of  weighing  to  1/500,000  mgm.,  and  found  to  be  about  222.4 
(density  of  oxygen  =  32). 

The  end-product  of  the  disintegration  is  lead,  and  all  uranium 
ores  contain  lead.  Lead  from  other  sources  gives  a  chloride  PbCl2 
in  which  207.20  parts  of  lead  are  combined  with  2  X  35.46  parts 
of  chlorine.  The  atomic  weight  207.2  cannot,  however,  be  reached 
by  subtracting  a  whole  number  of  atomic  weights  of  helium  from  the 
atomic  weight  of  uranium,  the  number  206  being  obtained  instead. 
Recently,  lead  chloride  prepared  from  the  lead  found  in  various  ores 
of  uranium  has  been  analyzed  by  Richards  of  Harvard,  as  well  as, 
independently,  by  two  other  chemists,  and  the  atomic  weight  of  this 
lead  was  found  to  be  from  206.1  to  206.8  in  different  samples.  This 
lead  chloride  has  properties  identical  with  those  of  ordinary  lead 
chloride  and  is,  therefore,  by  definition,  the  same  substance.  Hence 
these  investigations  have  revealed  the  first  known  exception  to  the 
law  of  definite  proportions.  Metallic  lead  from  radium  ores  has  the 
density  11.288  (Richards,  1916),  however,  that  of  ordinary  lead  being 
11.337  (19.94°),  which  gives  the  same  atomic  volume  for  both.  The 
spectra  of  these  leads  are  identical. 

Since  the  initial  (U)  and  final  (Pb)  materials  are  both  electrically 
neutral,  it  must  be  assumed  that  at  some  stages  more  than  one  electron 
per  atom  is  expelled.  8He++  are  lost  and  therefore  16e~. 

Additional  Data.  —  The  yield  of  radium  is  very  small.  6000 
kg.  of  pitchblende,  after  extraction  of  the  uranium,  yield  about 


872  INORGANIC  CHEMISTRY 

2000  kg.  of  residue.  This  affords  about  6  to  8  kg.  of  the  mixture 
of  radium  and  barium  sulphates,  from  which  0.2  g.  of  pure  radium 
bromide  can  be  prepared. 

One  gram  of  uranium,  after  it  has  produced  the  equilibrium 
proportion  of  radium  (about  3.2  X  10~7  g.),  gives  off  helium  at 
the  rate  of  1  c.c.  in  sixteen  million  years.  Since  the  mineral  fergu- 
sonite  contains  26  c.c.  of  accumulated  helium  for  every  gram  of 
uranium,  the  samples  of  this  mineral  must  be  at  least  416  million 
years  old. 

The  complete  disintegration  of  1  c.c.  of  niton  to  lead  would 
deliver  about  seven  million  calories,  but,  of  course,  the  liberation 
of  the  heat  would  be  spread  over  a  great  length  of  time. 

Chemical  Actions  of  the  "Rays."  —  The  radiations  which 
are  most  active  in  ionizing  air  and  in  acting  upon  photographic 
plates  are  the  ct-particles.  These  particles  also  cause  the  flashes 
of  light  when  they  encounter  zinc  sulphide.  The  radiations  change 
the  colors  of  minerals,  including  gems,  and  give  a  deep  violet  color 
to  the  glass  tube  containing  the  specimen.  They  also  turn  atmos- 
pheric oxygen  in  part  into  ozone  and,  in  solution,  produce  traces  of 
hydrogen  peroxide  in  the  water. 

The  radiations  also  destroy  minute  organisms  and  kill  the  cells 
of  the  skin,  producing  sores.  They  have  been  employed  in  the 
treatment  of  lupus  and  of  superficial  cancerous  growths. 

Other  Radioactive  Series.  —  Thorium,  found  as  phosphate  in 
monazite  sand,  is  also  radioactive  and  furnishes  a  series  of  disinte- 
gration products.  The  final  material  is  a  salt  of  lead.  Analysis 
of  the  chloride  of  lead  made  from  traces  of  the  element  found  in  all 
thorium  minerals  shows  that  the  atomic  weight  (Soddy)  is  208.4, 
while  that  of  ordinary  lead  is  207.2.  The  atom  of  thorium  (at.  wt. 
232.4)  thus  loses  6He  ( =  6  X  4  =  24)  during  the  disintegration. 
There  are  thus  three  chlorides  of  lead  with  identical  properties,  but 
different  compositions,  namely,  the  common  one  207.2  :  2X35.46, 
that  from  radium  206  :  2X35.46,  and  that  from  thorium  208.4  : 
2X35.46. 

Actinium  and  polonium  are  also  radioactive  elements,  which 
have  not  yet  been  fully  investigated.  The  former  appears  to  be 
formed  by  a  second,  parallel,  disintegration  of  Ui,  and  the  latter 
in  a  similar  way  from  Ra-E.  Compounds  of  potassium  and  rubidium 
show  traces  of  radioactivity. 


THE  CHROMIUM   FAMILY.      RADIUM  873 

Significance  of  Radioactivity.  —  The  Brownian  movement 
(p.  622)  has  revealed  to  us  bodies  intermediate  between  ordinary- 
particles  and  single  molecules,  and  has  enabled  us  to  estimate  the 
actual  weight  of  molecules.  Radioactivity  enables  us  to  count 
charged  molecules  of  helium  as  they  enter  the  electroscope  or  produce 
flashes  of  light  on  zinc  sulphide,  and  the  fog-tracks  permit  us  to 
follow  their  movements.  There  is  thus  now  no  question  that  mole- 
cules and  atoms  are  real.  Furthermore,  we  infer  that  all  kinds  of 
atoms  are  composed  of  a  positive  nucleus  (p.  470)  surrounded  by 
electrons,  although  only  the  atoms  of  radioactive  elements  are  un- 
stable. The  diameter  of  the  positive  nucleus  of  a  hydrogen  atom  is 
calculated  to  be  about  T-TOTF  °f  ^nat  of  an  electron.  Rutherford  has 
confirmed  this  by  actual  measurement.  The  atom  is  thus  no  longer 
regarded  as  being  solid  and  continuous  in  structure.  It  is  mainly  a 
vacuum,  containing  a  few  relatively  very  minute  bodies  possessing 
weight.  The  fact  that  a-particles  are  thus  able  to  plough  their  way 
through  molecules  of  oxygen  and  nitrogen,  being  diverted  from  a 
straight  path  only  when  they  happen  to  pass  very  close  to  the  positive 
nucleus  (which,  of  course,  repels  the  positive  a-particles),  is  no 
longer  mysterious. 

Another  interesting  conclusion  has  been  reached  from  the  ob- 
servation that  niton  is  found  in  the  soil  and  in  many  natural  waters. 
Calculation  shows  that  the  heat  given  off  by  the  disintegration  of 
the  amounts  of  radioactive  matter  known  to  exist  in  the  crust  of  the 
earth  is  alone  sufficient  to  account  for  the  maintenance  of  the  tem- 
perature of  the  planet.  A  globe  of  the  size  and  material  of  the  earth, 
possessing  originally  only  heat  energy,  and  cooling  from  a  white  hot 
condition  to  the  temperature  of  interstellar  space,  would  have  passed 
through  the  stage  of  habitable  temperatures  in  a  much  shorter  time 
than  that  which  a  study  of  the  geological  deposits  (and  the  fossils 
they  contain)  show  to  have  been  actually  available.  The  discovery 
of  the  enormous,  but  gradually  released  disintegration  energy  of  the 
radioactive  elements  enables  us  now  to  explain  the  prolonged  period 
during  which  life  has  existed  on  the  earth. 

Exercises.  —  1.  Construct  equations,  showing  the  interactions 
of:  (a)  chromic  oxide  and  aluminium,  (6)  strontium  nitrate  and 
potassium  dichromate  in  solution,  (c)  potassium  hydroxide  and 
chromic  hydroxide,  and  the  reversal  on  boiling,  (d)  chlorine  and 
potassium  chromite  in  excess  of  alkali  (what  is  the  actual  oxidizing 
agent?). 


874  INORGANIC   CHEMISTRY 

2.  What  volume  of  oxygen  at  0°  and  760  mm.,  (a)  is  obtainable 
from  one  formula-weight  of  potassium  dichromate  (par.  4,  p.  858), 
(6)  is  required  to  oxidize  one  formula-weight  of  chromous  chloride? 

3.  To  what  classes  of  actions  should  you  assign  the  three  methods 
of  making  chromic  oxide  (p.  861)? 

4.  Make  equations  for  all  the  reactions  involved  in  the  prepa- 
ration of  sodium-diuranate  from  pitchblende. 

5.  How   many   candle   power   will   be   obtained   from   50-watt 
carbon  and  tungsten  filament  lamps,  respectively? 

6.  Point  out  the  resemblance,  and  the  differences  between  the 
reactions  of,  (a)  gold  with  aqua  regia,  (b)  calcium  oxalate  with  hydro- 
chloric acid,  (c)  barium  chromate  with  nitric  acid  (p.  860). 


CHAPTER  XLIII 
MANGANESE 

The  Chemical  Relations  of  the  Element.  —  Manganese 
stands,  at  present,  alone  on  the  left  side  of  the  eighth  column  of  the 
periodic  table.  The  right  side  is  occupied  by  the  halogens.  It  is 
never  univalent,  as  are  the  halogens,  but  its  heptoxide  Mn2O7  and  the 
corresponding  acid,  permanganic  acid  HMn04,  are  in  many  ways 
closely  related  to  the  heptoxide  of  chlorine  and  perchloric  acid  HC1O4. 
Of  the  lower- oxides  of  manganese,  MnO  is  basic,  and  Mn2O3  feebly 
basic.  Mn02  is  feebly  acidic,  MnO3  more  strongly  so,  and  perman- 
ganic acid  (from  Mn2O?)  is  a  very  active  acid.  Contrary  to  the  habit 
of  feebly  acidic  and  feebly  basic  oxides,  such  as  those  of  zinc,  alumin- 
ium, and  tin,  the  basic  oxides  of  manganese  are  not  at  all  acidic,  and 
the  acidic  oxides,  with  the  exception  of  MnO2,  are  not  also  basic. 
There  are  thus  the  five  following,  rather  well-defined  sets  of  com- 
pounds, showing  five  different  valences  of  the  element.  Of  these  the 
first,  fourth,  and  fifth  are  the  most  stable  and  the  most  important. 

1.  Manganous  compounds,  MnO,  Mn(OH)2,  MnS04,  etc.      These 
compounds  resemble  those  of  the  magnesium  family  (and  those  of 
Fe"1"4").     The  salts  of  weak  acids,  such  as  the  carbonate  and  sulphide, 
are  easily  made,  and  there  is  little  hydrolysis  of  the  halides.     The 
salts  are  pale-pink  in  color. 

2.  Manganic  compounds,  Mn2O3,  Mn(OH)3,  Mn2(SO4)3,  [MnCl3]. 
The  salts  resemble  the  chromic  and  aluminium  salts  in  behavior,  but 
are  even  less  stable  than  are  those  of  quadrivalent  lead.     They  are 
completely  hydrolyzed  by  little  water.     The  salts  are  violet  in  color. 

3.  Manganites,   MnO2,  H2MnO3,  CaMnO3.     The  alkali  manga- 
nites  are  strongly  hydrolyzed,  like  the  plumbates  and  the  stannates. 

4.  Manganates,  Mn03,  H2MnO4,  K2MnO4.      The  salts  resemble 
the  sulphates  and  chromates,  but  are  much  more  easily  hydrolyzed. 
The  free  acid  resembles  chloric  acid  (p.  483)  in  that,  when  it  decom- 
poses, it  yields  a  higher  acid  (HMnO4)  and  a  lower  oxide  (Mn02). 
The  salts  are  green  in  color. 

5.  Permanganates,  Mn207,  HMn04  (hydrated),  KMn04.      The 
salts  resemble  the .  perchlorates,  and  are  not  hydrolyzed  by  water. 
They  are  reddish-purple  in  color. 

875 


876  INORGANIC   CHEMISTRY 

It  will  be  seen  that  the  element  manganese  changes  its  character 
totally  with  change  in  valence,  and  in  each  form  of  combination 
resembles  some  set  of  elements  of  valence  identical  with  that  which  it 
has  itself  assumed.  Since  the  valence  represents  the  number  of 
electrons  gained  or  lost  by  each  atom  (p.  793),  it  is  thus  evident 
that  the  chemical  properties  of  an  element  depend  more  upon  the 
electrical  constitution  of  its  atom  than  upon  the  atomic  weight. 
The  latter  is  a  secondary  property,  dependent  on  the  former  (cf. 
p.  470). 

Occurrence  and  Isolation.  —  The  chief  ore  is  the  dioxide, 
pyrolusite  MnO2,  which  always  contains  compounds  of  iron.  Other 
manganese  minerals  are:  braunite  Mn2O3;  the  hydrated  form,  man- 
ganiteMnO(OH);  hausmanniteMn3O4;  and  manganese  spar  MnC03. 
The  last  is  isomorphous  with  calcite.  The  metal  is  most  easily  made 
by  reducing  one  of  the  oxides  with  aluminium  by  Goldschmidt's 
method. 

Physical  and  Chemical  Properties.  —  The  metal  manganese 
(m.-p.  1260°)  has  a  grayish  luster  faintly  tinged  with  red.  It  is  oxi- 
dized superficially  by  air,  and  easily  displaces  hydrogen  from  dilute 
acids,  giving  manganous  salts.  Its  alloys  with  iron,  such  as  spiegel 
iron  (5-15  per  cent  Mn)  and  ferro-manganese  (70-80  per  cent  Mn), 
are  made  by  using  manganese  ores  with  the  charge  in  the  blast  fur- 
nace, and  are  added  to  the  iron  in  making  special  steels.  Manganese 
steel  (7-20  per  cent  Mn)  is  exceedingly  hard,  even  when  cooled  slowly. 
It  is  used  for  the  jaws  of  rock  crushing  machinery  and  for  burglar- 
proof  safes.  Wire,  made  of  an  alloy  called  manganin  (Cu  84  per 
cent,  Ni  4  per  cent,  Mn  12  per  cent),  invented  by  Weston,  is  used  in 
instruments  for  making  electrical  measurements,  because  its  resistance 
does  not  alter  with  moderate  changes  in  temperature. 

Oxides.  —  Manganous  oxide  MnO  is  a  green  powder,  made  by 
reducing  any  of  the  other  oxides  with  hydrogen.  Hausmannite 
Mn3O4  is  dull  red.  An  oxide  having  this  composition  is  formed  when 
any  of  the  other  oxides  is  heated  in  air,  oxidation  or  reduction,  as  the 
case  may  be,  taking  place  (cf.  p.  832).  This  oxide  corresponds  to 
minium  Pb3O4  (p.  832)  rather  than  to  Fe3O4,  for  with  dilute  acids 
it  gives  a  soluble  manganous  salt  and  a  precipitate  of  the  dioxide : 

Mn2MnO4  +  4HN03  -»  2Mn(N03)2  +  H4Mn04|. 


MANGANESE  877 

The  hydrated  dioxide  H4MnO4  subsequently  loses  water.  Haus- 
mannite  also  forms  square  prismatic  crystals.  In  view  of  its  behavior 
with  acids  and  its  crystalline  form,  it  is  thought  to  be  an  orthoman- 
ganite  of  manganese  Mn2MnO4,  rather  than  a  derivative  of  manganic 
oxide,  Mn(MnO2)2,  which  would  be  a  spinelle  (p.  812).  The  mag- 
netic oxide  of  iron  Fe(FeO2)2  belongs  to  the  regular  system,  like  the 
spinelles.  Manganic  oxide  Mn203  is  brownish-black,  and  is  formed 
by  heating  any  of  the  oxides  in  oxygen.  In  dilute  acids  it  behaves  as 
if  it  were  a  manganite  of  manganese  Mn.MnO3,  for  it  gives  a  man- 
ganous  salt  and  manganese  dioxide.  Yet  compounds  of  trivalent 
manganese  are  known,  and  this  may  be  one. 

Manganese  dioxide  MnO2  is  black,  and  is  most  easily  prepared  in 
pure  condition  by  gentle  ignition  of  manganous  nitrate.  The  hy- 
drated forms  of  the  oxide  are  produced  by  reactions  like  those  just 
mentioned,  and  by  adding  a  hypochlorite  or  hypobromite  to  man- 
ganous hydroxide  suspended  in  water.  Manganese  dioxide  is  not  a 
peroxidate  (p.  318).  That  is  to  say,  it  does  not  contain  the  radical 
(O2)  and,  therefore,  does  not  give  hydrogen  peroxide.  Its  reaction 
formula  is  Mn(O)2  not  Mn(O2)  and  in  double  decompositions  it 
yields  only  water  H2(O).  In  glass-making  (p.  726),  it  is  employed 
to  oxidize  the  green  ferrous  silicate,  derived  from  impurities  in  the 
sand,  to  the  pale-yellow  ferric  compound.  The  amethyst  color  of 
the  manganic  silicate  which  is  formed  tends  to  neutralize  this  yellow. 
The  dioxide  forms  the  depolarizer  ill  the  Leclanche"  cell  (p.  797). 
It  is  mixed  with  black  paints  as  a  " dryer"  (oxidizing  agent). 

Manganese  trioxide  MnO3  is  a  red,  unstable  powder.  Manganese 
heptoxide  Mn207  is  a  brownish-green,  volatile  oil  (see  below). 

When  any  of  these  oxides  is  heated  with  an  acid,  a  manganous  salt 
is  obtained.  Salts  of  this  class  are,  in  fact,  the  only  stable  substances 
in  which  manganese  is  combined  with  an  acid  radical.  In  this  action 
the  oxides  containing  more  oxygen  than  does  MnO  give  off  oxygen, 
or  oxidize  the  acid  (cf.  p.  219).  When  the  oxides  are  heated  with 
bases,  in  the  presence  of  air,  manganates  are  always  formed.  With 
the  oxides  containing  a  smaller  proportion  of  oxygen  than  MnO3, 
oxygen  is  taken  from  the  air. 

Manganous  Compounds.  —  The  manganous  salts  are  formed 
by  the  action  of  acids  upon  the  carbonate  or  any  of  the  oxides.  Thus 
the  chloride  MnCl2,4H20  is  obtained  in  pale-pink  crystals  from  a 
solution  made  by  treating  the  dioxide  with  hydrochloric  acid  and 
driving  off  the  chlorine  liberated  by  oxidation  (p.  219),  The  by- 


878  INORGANIC  CHEMISTRY 

droxide  Mn(OH)2  is  formed  as  a  white  precipitate  when  a  soluble 
base  is  added  to  a  solution  of  a  manganous  salt.  This  body  passes 
into  solution  when  ammonium  salts  are  added,  and  cannot  be  pre- 
cipitated in  their  presence  on  account  of  the  formation  of  molecular 
ammonium  hydroxide  and  the  suppression  of  the  hydroxide-ion 
(cf.  magnesium  hydroxide,  p.  766).  The  hydroxide  quickly  darkens 
when  exposed  to  the  air  and  passes  over  into  hydrated  manganic 
oxide  MnO(OH). 

Manganous  sulphate  gives  pink  crystals  of  a  hydrate.  Below  6° 
the  solution  deposits  MnSO4,7H2O,  which  is  a  vitriol  (p.  771).  Be- 
tween 7°  and  20°  the  product  is  MnSO4,5H2O,  asymmetric  and  isomor- 
phous  with  CuSO4,5H2O.  Above  25°  monosymmetric  prisms  of 
MnSC>4,4H2O  are  obtained.  These  hydrates  have  different  aqueous 
tensions  and  may  be  formed  from  one  another  by  lowering  or  raising 
the  pressure  of  water  vapor  around  the  substance  (p.  153).  The 
significance  of  the  temperatures  proper  to  the  crystallization  of  each 
(cf.  pp.  688,  728,  744)  is  that  a  given  solid  hydrate  can  be  formed 
only  in  a  solution  which  is  saturated  with  respect  to  that  hydrate 
and  has  the  same  aqueous  tension  as  the  hydrate.  These  conditions 
are  necessary  to  that  state  of  equilibrium  between  the  solution  and 
the  hydrate  on  which  the  co-existence  of  solution  and  hydrate  during 
crystallization  depends  (cf.  p.  195).  Hence  the  hydrates  with  the 
larger  proportions  of  water,  and  the  higher  aqueous  tensions,  are 
formed  in  the  colder  solutions  which  contain  less  of  the  solute  when 
saturated  and  have  therefore  at  a  given  temperature  themselves 
relatively  high  aqueous  tensions. 

The  presence  of  a  foreign  dissolved  body,  since  it  will  lower  the 
vapor  tension  of  the  solution,  may  similarly  cause  the  formation  of  a 
lower  hydrate.  Thus,  at  the  ordinary  temperature,  calcium  sulphate 
solution  has  a  higher  aqueous  tension  than  gypsum,  and  therefore 
gypsum  is  deposited  from  it,  and  anhydrite  will  turn  into  gypsum  if 
placed  in  it.  But  calcium  sulphate  solution  containing  much  of  the 
chlorides  of  sodium  and  magnesium  has  a  lower  aqueous  tension  than 
gypsum,  and  so  anhydrite  is  deposited,  and  gypsum  in  contact  with 
such  a  solution  would  lose  its  water  of  hydration.  This  explains  the 
deposition  of  anhydrite  in  the  salt  layers  (cf.  p.  717). 

Manganous  carbonate  MnCO3  is  a  white  powder  formed  by  pre- 
cipitation. The  sulphide  MnS  is  obtained  as  a  green,  crystalline 
powder  by  leading  hydrogen  sulphide  over  any  of  the  oxides.  A 
flesh-colored,  amorphous  variety  MnS  (often  somewhat  hydrated)  is 
more  familiar  and  is  precipitated  by  ammonium  sulphide  from  man- 


MANGANESE  879 

ganous  salts.  It  interacts  with  mineral  acids  and  even  with  acetic 
acid,  so  that  it  cannot  be  precipitated  by  hydrogen  sulphide  (cf. 
p.  774).  When  rubbed  in  a  mortar  it  becomes  crystalline,  and  is 
then  green. 

The  manganous  salts  of  weak  acids,  such  as  the  carbonate  and 
sulphide,  darken  when  exposed  to  air  and  are  oxidized,  with  formation 
of  hydrated  manganic  oxide.  As  we  have  seen,  manganousl  hydroxide 
is  similarly  oxidized  and  these  salts  are  precisely  the  ones  which 
should  furnish  the  hydroxide  by  hydrolysis.  While  there  is  a  general 
resemblance  between  the  manganous  salts  and  the  stannous,  chro- 
mous,  and  ferrous  salts,  the  manganous  salts  of  active  acids  are  not 
oxidized  by  the  air  as  are  the  corresponding  salts  of  the  other  three 
metals. 

Manganic  Compounds.  —  The  base  of  this  set  of  compounds, 
manganic  hydroxide  Mn(OH)3,  is  slowly  deposited  by  the  action  of 
the  air  on  an  ammoniacal  solution  of  a  manganous  salt  in  salts  of 
ammonium.  The  chloride  MnCl3  is  present  in  the  liquid  obtained 
by  the  action  of  hydrochloric  acid  upon  manganese  dioxide  (cf.  p.  219), 
but  loses  chlorine  very  readily  and  cannot  be  isolated.  Double  salts, 
however,  such  as  MnCl3,2KCl  and  MnF3,2KF,2H20,  are  known. 
Manganic  sulphate  Mn2(SO4)3  is  deposited  as  a  violet-red  powder 
when  hydrated  manganese  dioxide  is  heated  with  concentrated  sul- 
phuric acid  at  160°.  It  is  deliquescent  and  is  rapidly  hydrolyzed  in 
the  cold  even  by  a  little  water,  giving  the  brownish-black  hydroxide: 

Mn2(S04)3  +  6H20  r»  2Mn(OH)3  +  3H2SO4. 

The  caesium-manganic  alum  Cs2S04,Mn2(SO4)3,24H2O  seems  to  be 
the  most  stable  derivative. 

Manganites.  —  Although  manganese  dioxide  interacts  when 
fused  with  potassium  hydroxide,  simple  salts  derived  from  H2MnO3 
(=  H2O,MnO2)  or  H4Mn04  (=  2H2O,MnO2)  are  not  formed.  The 
products  are  complex,  as  K2Mn5On.  Some  less  complex  manganites 
are  formed  in  the  Weldon  process  for  utilizing  the  manganous  chlo- 
ride, formerly  obtained  in  manufacturing  chlorine.  The  liquor  is 
mixed  with  slaked  lime,  and  air  is  blown  through  the  mass  of  calcium 
and  manganous  hydroxides  which  is  thus  obtained.  Black  man- 
ganites of  calcium,  such  as  CaMn03  (=  CaO,Mn02)  and  CaMn205 
(CaO,2Mn02)  are  thus  formed: 

Ca(OH)2  +  2Mn(OH)2  +  02  ->  CaMn206  +  3H20, 


880  INORGANIC   CHEMISTRY 

and  when  afterwards  treated  with  hydrochloric  acid  they  behave  like 
mixtures  of  manganese  dioxide  and  calcium  oxide.  As  we  have  seen 
(p.  877),  the  oxides  MnsC^  and  Mn2O3  may  be  manganites  of  man- 
ganese. 

Manganates.  —  When  one  of  the  oxides  of  manganese  is  fused 
with  potassium  carbonate  (or  other  alkali)  and  potassium  nitrate  (or 
other  oxidizing  agent)  a  green  mass  is  obtained.  The  green  aqueous 
extract  deposits  potassium  manganate  K2MnO4  in  rhombic  crystals, 
which  are  isomorphous  with  those  of  potassium  sulphate,  and  are 
almost  black: 

K2C03  +  Mn02  +  0  ->  K2Mn04  +  C02. 

The  acid  H2MnO4,  itself  unknown,  must  be  weak,  for  the  potassium 
salt  is  easily  hydrolyzed.  The  salt  remains  unchanged  in  solution 
only  in  presence  of  free  alkali,  the  hydroxide-ion  of  the  alkali  combin- 
ing with  and  suppressing  the  hydrogen-ion  of  the  water  whose  com- 
bination with  the  MnOr1  ion  constitutes  the  hydrolysis.  When  the 
concentration  of  the  hydroxide-ion  is  reduced  by  dilution,  or,  better 
still,  when  a  weak  acid  such  as  carbonic  acid  or  acetic  acid  is  used  to 
neutralize  it,  the  salt  is  hydrolyzed,  according  to  the  partial  equation : 

•     K2MnO4  +  2H20  ->  2KOH  (+  H2Mn04) .  (1) 

The  free  acid  immediately  changes  so  that  a  part  is  oxidized  to  per- 
manganic acid,  giving  a  purple-red  color  to  the  solution,  and  a  part  is 
reduced  to  manganese  dioxide,  giving  a  black  precipitate.  The  trans- 
formation is  similar  to  that  of  chloric  acid  (p.  483).  The  equation 
may  be  made  by  noting  that  manganic  acid  has  the  composition 
H20,MnOa  and  changes  so  as  to  yield  H2O,Mn2O7  and  Mn02.  Thus 
each  molecule  of  H2Mn04,  in  forming  a  molecule  of  Mn02,  yields 
one  unit  of  oxygen,  while  2(H2O,MnOs)  +  O  are  required  to  give 
H20,Mn207  +  H20: 

3(H20,Mn03)  ->  H20,  Mn207  +  Mn02  +  2H20 
or  (3H2Mn04)  ->  2HMn04  +  Mn02  +  2H20.  (2) 

In  consequence  of  the  presence  of  potassium  hydroxide  (equation  (1)) 
the  product  is  potassium  permanganate : 

2KOH  +  2HMnO4  -»  2KMn04  +  2H2O.  (3) 

Multiplying  equation  (1)  by  3,  omitting  the  manganic  acid,  and 


MANGANESE  881 

adding  the  three  partial  equations,  we  have  the  equation  for  the 
action  as  it  really  occurs: 

3K2MnO4  +  2H2O  -»  4KOH  +  2KMn04  +  MnO2. 

To  make  the  equation  by  the  method  of  positive  and  negative 
values  (p.  493),  we  note  that  in  K2MnO4  we  have  2K+  and  4O~  and 
therefore  MnW  to  secure  electrical  neutrality.  The  latter  becomes 
Mnffl1*  and  Mntt.  Arithmetically  3Mnffi  will  give  2Mnttt+  and 
IMntt.  Hence,  3K2MnO4  are  required,  and  2KMnO4  and  !MnO2 
produced.  In  terms  of  the  ions  the  equation  is  simpler: 

3Mn04=  +  2H+  ->  20H~  +  2MnO4~  +  MnO2. 

The  alkaline  solution  of  potassium  manganate  interacts  readily 
with  oxidizable  substances.  Thus  oxalic  acid  is  converted  into  car- 
bonic acid,  and  alcohol  into  acetic  acid.  The  details  of  the  change 
depend  upon  the  amount  of  free  alkali  present  and  the  nature  of  the 
product  of  oxidation.  Lower  oxides  of  manganese  such  as  Mn02  are 
usually  precipitated. 

Permanganates.  —  Potassium  permanganate  KMn04  is  made 
by  hydrolysis  of  the  manganate  as  shown  above,  and  is  obtained,  as 
purple  crystals  with  a  greenish  luster,  by  evaporation  of  the  solution. 
The  crystals  are  rhombic  prisms,  isomorphous  with  potassium  per- 
chlorate.  To  avoid  the  loss  of  manganese  thrown  down  as  dioxide, 
the  action  is  carried  out  commercially  by  passing  ozone  through  the 
solution  of  the  manganate : 

2K2Mn04  +  03  +  H20  ->  2KMnO4  +  O2  +  2KOH. 

Sodium  permanganate  NaMnO4  is  made  in  a  similar  manner.  Alu- 
minium permanganate  in  solution  is  sold  as  "  Condy's  disinfecting 
fluid."  This  liquid  owes  its  properties  to  the  oxidizing  power  of  the 
permanganic  acid,  formed  by  hydrolysis  of  the  salt.  Permanganic 
acid  is  a  very  active  acid,  that  is,  it  is  highly  ionized  in  aqueous 
solution.  A  solid  hydrate  of  the  acid  may  be  secured  in  reddish- 
brown  crystals  by  adding  sulphuric  acid  to  a  solution  of  barium 
permanganate  and  allowing  the  filtrate  to  evaporate: 

Ba(Mn04)2  +  H2S04  -f  zH20  *±  BaS04J  +  2HMn04,zH2O. 

This  hydrate  decomposes,  on  being  warmed  to  32°,  and  yields  oxygen 
arid  manganese  dioxide.  When  a  very  little  dry,  powdered  potassium 
permanganate  is  moistened  with  concentrated  sulphuric  acid,  brown- 


882  INORGANIC  CHEMISTRY 

ish-green,  oily  drops  of  permanganic  anhydride  (manganese  heptoxide) 
Mn207  are  formed.  This  compound  is  volatile,  giving  a  violet 
vapor,  and  is  apt  to  decompose  explosively  into  oxygen  and  manganese 
dioxide.  Its  oxidizing  power  is  such  that  combustibles  like  paper, 
ether,  and  illuminating-gas  are  set  on  fire  by  contact  with  it. 

Potassium  Permanganate  as  an   Oxidizing  Agent.  —  The 

actions  are  different  according  as  the  substance  is  employed  (1)  in 
alkaline,  (2)  in  acid,  or  in  neutral  solution. 

1.  When  an  alkali,  such  as  potassium  hydroxide,  is  added,  the 
action  by  which  the  permanganate  is  formed  is  reversed,  and  the  solu- 
tion becomes  green  from  the  production  of  the  manganate: 

4KMnO4  +  4KOH  -»  4K2Mn04  +  2H20  +  Oa, 
or  4Mn04~  +  40H~  ->  4MnO4=  +  2H2O  +  02. 

When  a  substance  capable  of  being  oxidized  is  present,  the  reduction 
proceeds  further  and  manganese  dioxide  is  precipitated.  Schemati- 
cally: Mn207  —  >  2Mn02  +  3O,  so  that  two  molecules  of  the  perman- 
ganate, in  alkaline  solution,  can  furnish  three  chemical  units  of 
oxygen  to  the  oxidizable  body. 

2.  With  an  acid,  the  amount  of  oxygen  available  is  greater,  for  the 
manganous  salt  of  the  acid  is  formed:  Mn207  —  >  2MnO  +  5O.     Thus 
when  sulphuric  acid  is  added  to  potassium  permanganate  solution, 
and  sulphur  dioxide  is  led  through  the  mixture,  we  have  : 

2KMnO4  +  3H2SO4  ->  K2S04  +  2MnS04  +  3H2O  (+  5O)   (1) 
(5O)  +  5H2SO3^5H2SO4  (2) 


2KMnO4  +  3H2SO4+  5H2SO3^K2S04  +  2MnSO4  +  3H2O  +  5H2SO4 

In  this  case,  since  sulphuric  acid  is  a  product,  the  preliminary  addi- 
tion of  the  acid  was  superfluous.  In  other  cases,  the  partial  equation 
(1),  showing  the  available  5O,  remains  the  same,  while  the  other  par- 
tial equation  varies  with  the  substance  being  oxidized.  Thus,  with 
hydrogen  sulphide  as  reducing  agent,  we  have: 

(0)  +  H2S  -»  H20  +  S     X  5  (2r) 

and  with  ferrous  sulphate,  we  get  ferric  sulphate  : 

2FeS04  +  H2S04  (+  O)  ->  Fe2(S04)3  +  H20     X  5         (2") 

As  before,  (2')  and  (2")  must  be  multiplied  throughout  by  five,  before 
summation  is  made  (see  also  p.  320). 


MANGANESE  883 

Since  the  manganous  salt  is  colorless,  the  quantity  of  a  ferrous 
salt,  or  of  hydrogen  peroxide  (p.  320)  in  a  sample  of  a  solution  may 
be  measured  by  titrating  (p.  390)  the  solution  with  a  standard 
solution  of  potassium  permanganate  until  the  color  ceases  to  be 
destroyed,  and  then  noting  the  volume  used.  For  iron,  the  standard 
solution  may  be  prepared  so  that  1  c.c.  will  oxidize  0.01  g.  of  ¥e++. 

3.  When  dry  potassium  permanganate  is  heated,  it  decomposes 
as  follows:  2KMnO4  ->  K2Mn04  +  MnO2  +  O2.  The  neutral  solu- 
tion resembles  that  of  potassium  dichromate  in  oxidizing  substances 
which  are  reducing  agents,  but  is  more  active.  Thus  when  the 
powdered  salt  is  moistened  with  glycerine,  the  mass  presently  bursts 
into  flame.  The  fingers  are  stained  brown,  by  permanganates,  re- 
ceiving a  deposit  of  manganese  dioxide,  in  consequence  of  the  reducing 
power  of  the  unstable  organic  substances  in  the  skin.  The  destruc- 
tion of  minute  organisms  by  Condy's  fluid  results  from  a  similar 
action. 

Analytical    Reactions    of   Manganese    Compounds.  —  The 

ions  commonly  encountered  are  manganous-ion  Mn++,  which  is  very 
pale-pink  in  color,  permanganate-ion  MnO4~,  which  is  purple,  and 
manganate-ion  MnO4=,  which  is  green.  The  manganous  compounds 
give  with  ammonium  sulphide  the  flesh-colored  sulphide  which  inter- 
acts with  acids.  Bases  give  the  white  hydroxide,  which  darkens  by 
oxidation,  and  interacts  with  salts  of  ammonium.  The  black,  hy- 
drated  dioxide  is  precipitated  by  hypochlorites.  All  compounds  of 
manganese  confer  upon  the  borax  bead  an  amethyst  color  (manganic 
borate)  which,  in  the  reducing  flame,  disappears  (manganous  borate). 
A  bead  of  sodium  carbonate  and  niter  becomes  green,  the  manganate 
being  formed. 

Exercises.  —  1.  Consider  the  valence  of  manganese  in  the 
oxides  Mn3O4  and  Mn2Os,  on  the  theory  that  they  are  manganites. 

2.  What  do  we  mean  by  saying  that,  (a)  chromous  chloride  is 
stable  (p.  148),  but  easily  oxidized  by  the  air,  (6)  permanganic  acid 
is  an  active  acid,  (c)  permanganic  acid  is  an  active  oxidizing  agent  in 
presence  of  excess  of  an  acid? 

3.  Formulate  the  oxidations  of  hydrogen  sulphide,  of  ferrous 
sulphate,  of  oxalic  acid  (to  carbon  dioxide),  and  of  nitrous  acid  (to 
nitric  acid)  by  potassium  permanganate  in  acid  solution.     In  doing 
so,  employ  the  several  methods  suggested  on  pp.  269,  493-497. 


CHAPTER  XLIV 
IRON,    COBALT,   NICKEL 

THE  elements  iron  (Fe,  at.  wt.  55.84),  cobalt  (Co,  at.  wt.  58.97), 
and  nickel  (Ni,  at.  wt.  58.68)  are  not  corresponding  members  of 
successive  periods,  like  the  families  hitherto  considered.  They  are 
neighboring  members  of  the  first  long  period,  lying  between  its  first 
and  second  octaves  (p.  461),  and  form  a  transition  group  between  the 
adjoining  elements  within  those  octaves.  Thus,  iron  forms  ferrates 
M2IFeVIO4  and  ferric  salts  FemCl3,  as  well  as  ferrous  salts  FenCl2, 
These  resemble  the  chromates  and  manganates,  the  chromic  and 
manganic  salts,  and  the  chromous  and  manganous  salts,  respectively. 
Cobalt  forms  cobaltic  and  cobaltous  salts,  like  Co2m(SO4)3  and  ConCl2. 
Nickel  enters  only  into  nickelous  salts,  like  NiCl2,  and  thus  links  iron 
and  cobalt  with  copper  and  zinc  which  are  both  bivalent  elements. 
The  free  metals  of  this  family  are  magnetic,  iron  showing  this  property 
strongly  and  cobalt  very  distinctly. 

IRON  FE 

Chemical  Relations  of  the  Element.  —  The  oxides  and  hy- 
droxides FeO  and  Fe(OH)2,  Fe2O3  and  Fe(OH)3  are  basic,  the  former 
more  strongly  so  than  the  latter.  The  ferrous  salts,  derived  from 
Fe(OH)2,  resemble  those  of  the  magnesium  group  and  those  of 
Cr++  and  Mn"1"1"  and  are  little  hydrolyzed.  The  ferric  salts,  derived 
from  Fe(OH)3,  resemble  those  of  Cr+++  and  Al+++  and  are  noticeably 
hydrolyzed.  Ferric  hydroxide  is  even  less  acidic,  however,  than  is 
chromic  hydroxide.  Iron  gives  also  a  few  ferrates  K2FeO4,  CaFe04, 
etc.,  derived  from  an  acid  H2FeO4  which,  like  manganic  acid  H2Mn04, 
(p.  880),  is  too  unstable  to  be  isolated.  Complex  anions  containing 
this  element,  such  as  the  anion  of  K4.Fe(CN)6,  are  familiar,  but  com- 
plex cations  containing  ammonia  are  unknown. 

The  ferrous  salts  differ  from  most  of  the  manganous  salts  and  re- 
semble the  chromous  and  stannous  salts  in  being  easily  (although 
not  quite  so  easily)  oxidized  by  the  air,  passing  into  the  ferric  con- 
dition. 

884 


IRON,   COBALT,   NICKEL  885 

Occurrence.  —  Free  iron  is  found  in  minute  particles  in  some 
basalts,  and  many  meteorites  are  composed  of  it.  Meteoric  iron  can 
be  distinguished  from  specimens  of  terrestrial  origin  by  the  fact  that 
it  contains  3-8  per  cent  of  nickel.  The  chief  ores  of  iron  are  the 
oxides;  haematite  Fe2Os  and  magnetite  FeaO^  and  the  carbonate 
FeCOs,  siderite.  The  first  is  reddish  and  columnar  in  structure;  but 
black,  shining,  rhombohedral  crystals,  known  as  specularite,  are  also 
found.  Hydrated  forms,  like  brown  iron  ore  2Fe2O3,3H20,  are  also 
common.  Siderite  is  pale-brown  in  color  and  rhombohedral,  isomor- 
phous  with  calcite.  When  mixed  with  clay  it  forms  iron-stone  and, 
with  20-25  per  cent  of  coal  in  addition,  black-band.  Most  of  the  iron 
in  Great  Britain,  but  less  than  one  per  cent  of  that  in  the  United 
States  is  obtained  from  these  two  sources.  Pyrite  FeS2  consists  of 
golden-yellow,  shining  cubes  or  pentagonal  dodecahedra.  It  is  used, 
on  account  of  its  sulphur,  in  the  manufacture  of  sulphuric  acid,  but, 
from  the  oxidized  residue,  iron  of  sufficient  purity  is  obtained  with 
difficulty.  Compounds  of  iron  are  associated  with  chlorophyll  and 
are  found  in  the  blood  (haemoglobin),  and  doubtless  play  an  impor- 
tant part  in  connection  with  the  vital  functions  of  these  substances. 
By  interaction  with  organic  compounds  of  iron  present  in  the  tissues, 
ammonium  sulphide  blackens  the  skin,  ferrous  sulphide  being  formed. 

Pure  Iron.  —  Pure  iron  is  obtained  by  reducing  pure  ferrous 
oxalate  in  a  stream  of  hydrogen  at  a  high  temperature.  It  is  also 
made  by  electrolysis  of  ferrous  sulphate  solution  at  100°  between 
iron  electrodes.  It  is  silver-white  and  melts  at  1530°.  The  purest 
iron  does  not  rust  in  pure  cold  water,  but  the  impurities  in  ordi- 
nary iron  act  as  contact  agents  and  rusting  proceeds. 

Metallurgy.  —  The  ores  of  iron  are  first  roasted  in  order  to  de- 
compose carbonates  and  oxidize  sulphides  if  these  salts  are  present. 
Coke  is  then  used  to  reduce  the  oxides.  Coal  is  unsuitable  because 
so  much  heat  is  wasted  in  driving  out  the  volatile  matter  and  moisture, 
which  are  absent  from  coke.  Ores  containing  lime  or  magnesia  are 
mixed  with  an  acid  flux,  such  as  sand  or  clay-slate,  in  order  that  a 
fusible  slag  may  be  formed.  Conversely,  ores  containing  silica  and 
clay  are  mixed  with  limestone.  With  proper  adjustment  of  the  in- 
gredients the  process  can  be  carried  on  continuously  in  a  blast  furnace 
(Fig.  164),  an  iron  structure  40  to  100  feet  high,  lined  with  firebrick. 
The  solid  materials  thrown  in  at  the  top  are  converted,  as  they  slowly 
descend,  completely  into  gases  which  escape  and  liquids  (iron  and 


886 


INORGANIC   CHEMISTRY 


FIG.  164. 


slag)  which  are  tapped  off  at  the  bottom.  Heated  air  is  blown  in  at 
the  bottom  through  tuyeres,  and  the  top  is  closed  by  a  cone  which 
descends  for  a  moment  when  an  addition  is  made  to  the  charge. 
The  gases,  which  contain  much  carbon  monoxide, 
are  led  off  and  used  to  heat  the  blast  or  to  drive 
gas-engines. 

The  main  action  takes  place  between  the  carbon 
monoxide,  present  in  consequence  of  the  excess  of 
carbon,  and  the  oxide  of  iron: 

Fe304  -f-  4CO  <=>  3Fe  +  4CO2. 
Since  the  action  is  a  reversible  one,  a  large  excess 
of  carbon  monoxide  is  required.  At  650°,  equilib- 
rium is  reached  with  CO  :  CO2  ::  1  vol.  :  1J  vols., 
at  800°,  1  vol.  :  13  vols.,  and  in  practice  the  pro- 
portion of  carbon  monoxide  used  is  from  twice  to 
fifteen  times  as  great.  Almost  5  tons  of  air,  heated 
in  advance  to  800°,  are  required  for  each  ton  of 
iron  produced.  The  moisture  in  this  air  acts  upon 
the  coke,  giving  water-gas  (p.  577).  This  action 
uses  up  fuel,  and  also  lowers  the  temperature  at 
the  point  where  it  should  be  highest.  In  some  plants  the  Gayley 
dry  blast  process  was  used  for  a  time,  but  careful  regulation  of  the 
blast,  in  accordance  with  the  humidity  of  the  air,  is  now  the  pre- 
vailing practice.  In  the  United  States  alone  31  million  tons  of  pig 
iron  are  annually  produced  (1913).  This  is  considerably  over  40 
per  cent  of  the  world's  production,  20  per  cent  being  supplied  by 
Germany  and  15  per  cent  by  Great  Britain. 

In  the  upper  part  of  the  furnace,  the  heat  (400°)  loosens  the 
texture  of  the  ore.  Further  down,  the  temperature  is  higher  (500- 
900°),  and  the  carbon  monoxide  reduces  the  oxide  of  iron  to  particles 
of  soft  iron.  A  temperature  high  enough  to  melt  pure  iron  is  barely 
reached  anywhere  in  the  furnace,  but,  a  little  lower  down,  by  union 
with  carbon,  the  more  fusible  cast  iron  (m.-p.  about  1200°)  is  formed 
and  falls  in  drops  to  the  bottom.  It  is  in  this  region  also  that  the  slag, 
essentially  a  glass,  is  produced.  If  the  flux  had  begun  sooner  to 
interact  with  the  unreduced  ore,  iron  would  have  been  lost  by  the 
formation  of  the  silicate.  The  iron  collects  below  the  slag,  and  the 
latter  flows  continuously  from  a  small  hole.  The  former  is  tapped 
off  at  intervals  of  six  hours  or  so  from  a  lower  opening.  As  a  rule, 
the  iron  never  cools  until  it  has  been  converted  into  rails  or  structural 
iron.  In  some  cases,  it  is  made  into  "pigs"  in  a  casting  machine. 


IRON,   COBALT,   NICKEL  887 

Cast  Iron  and  Wrought  Iron.  —  Pure  iron  is  not  manufac- 
tured, and  indeed  would  be  too  soft  for  most  purposes.  Piano-wire, 
however,  is  about  99.7  per  cent  pure.  The  product  obtained  from 
the  blast  furnace  contains  92-94  per  cent  of  iron  along  with  2.6-4.3 
per  cent  of  carbon,  often  nearly  as  much  silicon,  varying  proportions  of 
manganese,  and  some  phosphorus  and  sulphur.  The  last  four  in- 
gredients are  liberated  from  combination  with  oxygen  by  the  carbon 
in  the  hottest  part  of  the  furnace  and  combine  or  alloy  themselves 
with  the  iron.  Cast  iron  does  not  soften  before  melting,  as  does  the 
purer  wrought  iron,  but  melts  sharply  at  1150-1250°  according  to  the 
amount  of  foreign  material  it  contains.  When  suddenly  cooled  it 
gives  chilled  cast  iron  which  is  very  brittle  and  looks  homogeneous  to 
the  eye,  all  the  carbon  being  present  in  the  form  of  carbide  of  iron 
Fe3C  (cementite)  in  solid  solution  in  the  metal.  By  slower  cooling, 
time  is  permitted  for  the  separation  of  part  of  the  carbon  as  graphite, 
which  appears  in  tiny  black  scales  (see  below),  and  gray  cast  iron 
results.  This  mixture  is  much  softer,  on  account  of  the  amount  of 
free,  relatively  pure  iron  which  it  contains. 

Wrought  iron,  invented  by  Henry  Cort  (1784),  is  made  by  heating 
the  broken  "  pigs  "  of  cast  iron  upon  a  layer  of  material  containing 
oxide  of  iron  and  hammer-slag  (basic  silicate  of  iron)  spread  on  the 
bed  of  a  reverberatory  furnace  (Fig.  143,  p.  686).  The  carbon,  silicon, 
and  phosphorus  combine  with  the  oxygen  of  the  oxide,  and  the  last 
two  pass  into  the  slag.  The  sulphur  is  found  in  the  slag  as  ferrous 
sulphide.  On  account  of  the  effervescence  due  to  the  escape  of  car- 
bon monoxide,  the  process  is  called  "  pig-boiling."  The  iron  is  stirred 
with  iron  rods  ("  puddled  ")  and  stiffens  as  it  becomes  purer,  until 
finally  it  can  be  withdrawn  in  balls  ("  blooms  ")  and  be  partially 
freed  from  slag  by  rolling.  The  resulting  bars  are  repeatedly  cut, 
piled  in  a  bundle,  reheated,  and  rolled.  The  iron  now  softens  suf- 
ficiently for  welding  below  1000°  and  melts  at  1505°  or  lower,  accord- 
ing to  its  purity.  If  it  still  contains  more  than  a  trace  of  combined 
phosphorus  it  is  brittle  when  cold  ("  cold  short ").  A  little  surviving 
sulphide  of  iron  makes  it  brittle  when  hot  ("  red-short ")  and  un- 
suitable for  forging.  Wrought  iron  should  contain  only  0.1-0.2  per 
cent  of  carbon.  Its  fibrous  structure  is  due  partly  to  the  films  of 
slag  which  have  not  been  completely  pressed  out  by  the  rolling.  On 
account  of  its  toughness,  wrought  iron  is  used  for  anchors,  chains, 
and  bolts,  and  for  drawing  into  wire.  On  account  of  its  relative 
purity  (99.8-99.9  per  cent),  it  is  less  fusible  than  cast  iron  and  is  used 
for  fire  bars.  The  above  operations  are  now  performed  by  machinery, 


888  INORGANIC   CHEMISTRY 

but  have  been  largely  displaced  by  the  Bessemer  and  open  hearth 
processes  in  which  iron  of  equal  purity  can  be  obtained. 

Properties  of  Steel.  —  This  is  a  variety  of  iron  almost  free 
from  phosphorus,  sulphur,  and  silicon.  Tool-steel  contains  0.9-1.5 
per  cent  of  carbon,  structural  steel  only  0.2-0.6  per  cent,  and  mild 
steel  0.2  per  cent  or  even  less.  Steel  combines  the  properties  of 
cast  and  of  wrought  iron,  being  hard  and  elastic,  and  at  the  same 
time  available  for  forging  and  welding  when  the  proportion  of  carbon 
is  low.  Steel  can  be  tempered  (see  below).  It  has  also  a  greater 
tensile  strength*  than  has  wrought  iron,  and  it  can  be  permanently 
magnetized. 

Bessemer  Process.  —  Steel  is  made  largely  by  the  Bessemer 
process  (Kelly  1852,  Bessemer  1855).  The  molten  cast  iron  is 
poured  into  a  converter  (Fig.  165)  and  a  blast  of  air  (a)  is  blown 

through  it.  The  oxidation  of  the 
manganese,  carbon,  silicon,  and  a 
little  of  the  iron  gives  out  sufficient 
heat  to  raise  the  temperature  of  the 
mass  above  the  melting-point  of 
wrought  iron.  The  required  propor- 
tion of  carbon  is  then  introduced  by 
adding  pure  cast  iron,  spiegel  iron, 
or  coke,  and  the  contents,  first  the 
slag,  and  then  the  molten  steel,  are 

finally  poured  out  by  turning  the  converter.  When  the  cast  iron 
contains  much  phosphorus,  the  oxide  of  this  element  is  reduced 
again  by  the  iron  as  fast  as  it  is  formed  by  the  blast.  In  such 
cases  a  basic  lining  containing  lime  and  magnesia  takes  the  place  of 
the  sand  and  clay  lining  of  the  ordinary  Bessemer  converter,  and  a 
slag  containing  a  basic  phosphate  of  calcium  is  produced.  This 
modification  constitutes  what  is  known  as  the  Thomas- Gilchrist 
process.  The  slag  (Thomas-slag)  when  pulverized  forms  a  valuable 
fertilizer  (cf.  p.  720).  In  the  United  States,  the  basic  open-hearth 
process  is  preferred. 

Being  obtained  from  molten  material,  steel  and  cast  iron  are  free 
from  slag. 

*  Tensile  strength  or  tenacity  is  measured  by  the  weight  (in  kilos)  required 
to  break  a  wire  of  the  metal  1  sq.  mm.  in  section.  Lead  2.6,  copper  51,  iron  71, 
steel  91. 


IRON,   COBALT,   NICKEL 


889 


Open-Hearth  (Siemens- Mar  tin)  Process.  —  In  this  process 
the  cast  iron  is  melted  in  a  saucer-shaped  depression  (Fig.  166), 
which  is  lined  with  sand  in  the  acid  process,  and  with  lime  and  mag- 
nesia in  the  basic  process.  Scraps  of  iron  plate  (for  dilution)  and 
haematite,  or  some  other  oxide  ore,  are  then  added  in  proper  pro- 
portions. The  materials  (50-75  tons  in  one  charge)  are  heated  with 
gas  fuel  for  8-10  hours.  To  secure  economically  the  high  temperature 
required  to  keep  the  product  (almost  pure  iron)  fused,  Siemens  devised 
the  method  of  preheating  the  fuel  gas  and  air  by  a  regenerative  device. 


FIG.  166. 


The  spent  air  and  gas  pass  down  through  a  checkerwork  of  brick. 
When  this  becomes  heated,  the  valves  are  reversed,  the  gas  and  air 
now  enter  through  the  heated  brickwork  and,  after  meeting  and 
burning  over  the  iron,  pass  out  through  the  checkerwork  on  the  op- 
posite side,  raising  its  temperature  in  turn. 

The  changes  are  similar  to  those  in  the  Bessemer  process.  During 
casting,  some  aluminium  is  added  to  combine  with  oxygen  (present 
as  CO)  and  give  sounder  ingots.  Recently,  iron  containing  10-15  per 
cent  of  titanium  has  been  added  instead.  The  titanium  combines 
with  both  nitrogen  and  oxygen  and  the  compounds  pass  into  the  slag, 
just  as  does  aluminium  oxide.  Rails  made  of  steel  purified  with  thi& 
element  are  less  liable  to  breakage  (the  commonest  cause  of  wrecks) 
and  are  40  per  cent  more  durable  than  are  ordinary  open-hearth  rails. 


890  INORGANIC  CHEMISTRY 

The  advantage  of  the  open-hearth  process  over  that  of  Bessemer 
is  that  it  is  not  hurried,  and  is  therefore  under  better  control.  The 
material  can  be  tested  by  sample  at  intervals  until  the  required 
composition  has  been  reached.  The  produce  is  of  more  uniform 
quality.  When  fine  steel  is  required,  electric  heating  (e.g.,  in  the 
Heroult  furnace)  permits  even  more  deliberate  treatment. 

Bessemer  and  open-hearth  steels  are  used  for  heavy  and  light 
machinery  castings  and  for  shafts.  It  is  rolled  into  rails,  and  into 
bridge  and  structural  iron. 

Crucible  Steel.  —  For  special  purposes  steel  is  made  in  cru- 
cibles of  clay  (or  graphite  and  clay)  in  melts  of  60-100  pounds. 
"  Melting  bar,"  a  very  pure  open-hearth  steel,  is  melted  with  charcoal 
or  with  pure  pig  iron.  This  steel  is  employed  in  making  razors 
(1.5  per  cent  C),  tools  (1  per  cent  C),  dies  (0.75  per  cent  C),  pens, 
needles,  and  cutlery.  Special  steels,  containing  one  or  more  of  the 
elements  manganese,  nickel,  chromium,  tungsten,  molybdenum,  and 
vanadium  are  made  in  this  way. 

Tempering.  —  When  steel  is  heated  to  redness  and  cooled  slowly, 
it  is  comparatively  soft.  Sudden  chilling,  however,  renders  it  harder 
than  glass.  By  subsequent,  cautious  heating  the  hardness  may  be 
reduced  to  any  required  extent,  and  this  treatment  is  called  tempering. 
The  sufficiency  of  the  heating  is  judged  roughly  by  the  interference 
colors  caused  by  the  thin  film  of  oxide  which  forms  on  the  surface. 
Thus  a  pale-yellow  color  (200-220°)  serves  for  tempering  razors,  a 
decided  yellow  (240°)  for  pen-knives,  a  brown  (260°)  for  shears, 
a  purple  (275°)  for  table-knives,  a  blue  (300-315°)  for  watch-springs 
and  sword-blades,  and  a  black-blue  (340°)  for  saws.  Except  in  the 
case  of  watch-springs,  these  films  are  afterwards  removed  by  the 
grinding. 

To  understand  this  behavior  it  must  be  noted  that  there  are  three 
states  of  solid  iron  resembling  the  rhombic  and  monoclinic  states 
of  sulphur  (cf.  p.  412).  The  form  stable  below  760°  is  known  as 
a-f errite  (wrought  iron) ,  It  is  magnetic  and  can  hold  little  carbide 
of  iron  in  solid  solution.  Above  760°  this  changes,  with  absorption  of 
heat,  into  /3-ferrite  which,  likewise,  holds  little  of  the  carbide  in 
solution,  but  is  not  magnetic.  At  900°  this  changes,  with  further 
absorption  of  heat,  into  7-ferrite,  a  non-magnetic  form  in  which  the 
carbide  is  soluble.  When  allowed  to  cool,  iron  assumes  these  forms 
in  the  reverse  order.  If,  now,  a  fluid  solution  of  carbon  in  iron, 


IRON,   COBALT,   NICKEL  891 

suitable  for  steel,  is  suddenly  chilled,  a  great  part  of  the  cold  mass  is  a 
supercooled  solid  solution  of  carbon  in  7-ferrite.  This  solid  solution 
is  called  martensite  and  is  very  hard  and  brittle.  It  is  less  stable 
at  ordinary  temperatures  than  is  a-ferrite,  but,  as  is  the  case  with 
yellow  phosphorus  (p.  551)  and  amorphous  sulphur  (p.  413),  the  low 
temperature  having  once  been  reached,  transformation  into  the  more 
stable  form  is  thereafter  exceedingly  slow.  The  material  is  hard  steel. 

When  the  molten  steel  (solution  of  carbon  in  iron)  is  allowed  to 
cool  so  slowly  that  equilibrium  can  be  reached  at  every  step,  a  compli- 
cated series  of  changes  ensues.  First  the  mass  solidifies  (at  or  before 
1130°)  to  a  mixture  of  martensite  (7-ferrite  with  carbon  in  solid  solu- 
tion up  to  2  per  cent)  and  graphite.*  As  the  temperature  now  falls 
very  slowly,  more  graphite  separates  until,  at  1000°,  1.8  per  cent 
remains  in  solution.  From  this  point  the  dissolved  carbide  of  iron 
(cementite  Fe3C  containing  6.6  per  cent  of  carbon)  is  separated.  At 
670°  pure  a-ferrite  also  begins  to  appear  and  there  remains  only  about 
0.9  per  cent  carbon  in  solid  solution.  At  this  temperature,  if  sufficient 
time  is  allowed,  the  solid  solution  separates  into  a  mixture  of  pure 
iron  (87  per  cent)  which  is  soft  and  carbide  of  iron  (13  per  cent)  which 
is  hard.  The  final  result  is  a  mechanical  mixture  of  a-ferrite  (wrought 
iron),  carbide  of  iron,  and,  if  the  original  amount  of  carbon  was 
sufficiently  large,  graphite.  These  components  may  be  recognized 
by  making  a  microscopic  study  of  a  polished  surface,  and  their 
formation  may  be  followed  by  chilling  the  specimen  at  any  desired 
stage.  The  soft  iron  which  predominates  in  the  product  of  slow 
cooling  makes  the  whole  soft.  Heating  to  a  high  temperature  and 
sudden  chilling  gives  the  homogeneous  solid  solution  of  carbon  in 
7-ferrite  once  more  and  restores  the  qualities  characteristic  of  steel. 
Moderated  reheating  (tempering)  of  the  chilled  mass  results  in  more 
or  less  partial  accomplishment  of  the  changes  proper  to  slow  cooling, 
and  consequently  in  a  more  or  less  close  approach  to  the  condition 
which  results  from  this. 

The  difference  between  the  effect  of  rapid  and  slow  cooling  of  cast 
iron  (p.  887)  can  now  be  made  clear.  Rapid  cooling  leads  to  the 
omission  of  the  intervening  steps  enumerated  above  and,  if  some- 
thing like  5  or  6  per  cent  of  carbon  is  present,  the  material  turns 
almost  completely  into  martensite.  This  is  chilled  cast  iron.  With 
slower  cooling,  much  graphite  separates,  and  the  product,  gray 
•cast  iron,  contains  much  less  of  the  carbide  and  much  more  free  iron. 

*  When  molten  cast  iron,  containing  3-4.5  per  cent  of  carbon,  is  cooled  in  this 
fashion,  the  amount  of  graphite  may  be  considerable. 


892  INORGANIC  CHEMISTRY 

The  various  changes  which  occur  in  cooling  steel  are  retarded  by 
the  presence  of  foreign  substances,  just  as,  with  sulphur  (p.  413), 
foreign  substances  delay  the  change  from  SM  to  Sx  and  permit  the 
supercooling  of  the  former  and  its  appearance  in  the  form  of  amor- 
phous sulphur.  Manganese,  nickel,  and  other  metals,  in  particular, 
greatly  reduce  the  facility  with  which  7-ferrite  passes  into  0-  and 
a-ferrite  at  900°  and  760°.  Thus  iron  with  12  per  cent  of  manganese, 
when  chilled  from  a  high  temperature,  contains  only  supercooled 
7-ferrite  and  is  non-magnetic.  It  has  to  be  kept  for  hours  (instead  of 
a  few  minutes)  at  a  temperature  below  760°,  say  500-600°,  before  it 
goes  over  into  a-ferrite.  Manganese  is  thus  a  valuable  constituent  of 
steel  because,  by  favoring  the  survival  of  the  7-ferrite  in  which  alone 
the  carbon  is  soluble,  it  permits  the  manufacture  of  a  homogeneous 
steel  containing  an  unusually  large  proportion  of  dissolved  carbon, 
and  allows  slower  cooling  without  loss  of  temper  (c/.  p.  890). 

Steel  Alloys.  —  As  we  have  seen,  substances  such  as  alumin- 
ium, titanium,  and  ferrosilicon  are  added  to  iron  for  the  purpose  of 
purifying  it,  and  pass  in  combination  into  the  slag.  There  are, 
however,  regular  alloys  containing  the  foreign  metal  along  with  the 
iron.  Thus,  manganese  steel  (7-20  per  cent  Mn),  made  by  adding 
spiegel  iron  or  ferromanganese  (p.  876)  to  steel,  remains  hard  even 
when  cooled  slowly  and  is  used  for  the  jaws  of  rock-crushers  and  for 
safes.  Chromium- vanadium  steel  (1  per  cent  Cr,  0.15  per  cent  Va) 
has  great  tensile  strength,  can  be  bent  double  while  cold,  and  offers 
great  resistance  to  changes  of  stress  and  to  torsion.  It  is  used  for 
frames  and  axles  of  automobiles  and  for  connecting  rods.  Tungsten 
steel  has  already  been  described  (p.  864).  Nickel  steel  (2-4  per  cent 
Ni)  resists  corrosion,  has  a  high  limit  of  elasticity  and  great  hardness, 
and  is  used  for  armor-plate,  wire  cables,  and  propeller  shafts.  Invar 
(36  per  cent  Ni)  is  practically  non-expansive  when  heated  within 
narrow  limits  and  is  used  for  meter-scales  and  pendulum  rods. 

Chemical  Properties  of  Iron.  —  Although  the  purest  iron  does 
not  rust  in  cold  water  (p.  885),  ordinary  iron  rusts  in  moist  air  or 
under  water.  It  probably  rusts  in  water  free  from  carbon  dioxide, 
displacing  the  hydrogen-ion,  but  the  action  is  greatly  hastened  by 
the  presence  of  carbonic  acid.  Rust  is  a  brittle,  porous,  loosely 
adherent  coating  of  variable  composition,  consisting  mainly  of  A 
hydrated  ferric  oxide  3Fe2O3,H2O,  which  does  not  protect  the  metal 
below.  Oil  protects  iron  from  rusting  because,  although  oxygen 


IRON,   COBALT,   NICKEL  893 

is  more  soluble  in  most  oils  than  in  water  (p.  87),  and  so  reaches  the 
iron  freely,  water  is  not  soluble  in  oil  and  so  moisture  is  excluded. 

Iron  burns  in  oxygen  and  interacts  with  superheated  steam,  in 
both  cases  giving  Fe304.  A  superficial  layer  of  this  oxide  adheres 
firmly  and  protects  the  iron  from  the  action  of  the  air  (Barff's  process 
iron,  or  Russia  iron). 

Iron  displaces  hydrogen  easily  from  dilute  acids.  Steel  and  cast 
iron,  which  contain  iron,  its  carbide,  and  graphite,  give  with  cold 
dilute  acids  almost  pure  hydrogen,  and  the  carbide  and  graphite 
remain  unat tacked.  More  concentrated  acids,  however,  particularly 
when  warm,  give  off,  along  with  hydrogen,  hydrocarbons  formed  by 
interaction  with  the  carbide  (p.  589).  The  odor  of  the  gas  is  due  to 
compounds  of  sulphur  and  phosphorus.  With  dilute  nitric  acid,  iron 
gives  ferrous  nitrate  and  ammonium  nitrate  (c/.  tin,  p.  824)  and  with 
the  concentrated  nitric  acid  ferric  nitrate  and  oxides  of  nitrogen. 
Iron  has  little  action  upon  alkalies. 

Although  iron  acts  vigorously  on  dilute  or  concentrated  nitric 
acid,  it  is  indifferent  to  fuming  nitric  acid  (NO2  in  solution,  p.  527). 
It  becomes  passive.  In  this  state,  it  no  longer  displaces  hydrogen 
from  dilute  acids.  If  dipped  in  cupric  sulphate  solution,  it  does 
not  receive  the  usual  red  coating  of  metallic  copper.  However,  if 
scratched  or  struck,  the  passive  condition  is  destroyed,  and  copper 
begins  to  be  deposited  at  the  point  touched  and  the  action  spreads 
quickly  over  the  whole  surface.  No  satisfactory  explanation  of  this 
phenomenon  has  been  obtained,  although  it  is  shown  also  by  chro- 
mium (p.  855),  cobalt,  and  other  metals. 

Ferrous  Compounds.  —  Ferrous  chloride  is  obtained  as  a  pale- 
blue  hydrate  FeCl2,4H2O  (turning  green  in  the  air)  by  interaction  of 
hydrochloric  acid  with  the  metal  or  the  carbonate.  The  anhydrous 
salt  sublimes  in  colorless  crystals  when  hydrogen  chloride  is  led  over 
the  heated  metal.  At  a  high  temperature  the  vapor  of  ferrous 
chloride  has  a  density  corresponding  to  the  simple  formula  FeCl2, 
but  at  lower  temperatures  there  is  much  association  (p.  282)  and  the 
formula  approaches  Fe2Cl4.  In  solution  the  salt  is  oxidized  by  the 
air  to  a  basic  ferric  chloride : 

4Fe++  +  O2  +  2H2O  ->  4Fe+++  +  40H~. 

In  presence  of  excess  of  the  acid,  normal  ferric  chloride  is  formed. 

With  nitric  acid,  ferric  chloride  and  nitric  oxide  are  produced  (p.  535). 

Ferrous  hydroxide  Fe(OH)2  is  thrown  down  as  a  white  precipitate, 


894  INORGANIC   CHEMISTRY 

in  water  freed  from  dissolved  oxygen  by  boiling,  but  rapidly  be- 
comes dirty-green  and  finally  brown,  by  oxidation.  It  interacts 
with,  and  dissolves  in  solutions  of  salts  of  ammonium,  being,  like 
magnesium  hydroxide  (p.  766),  sufficiently  soluble  in  water  to  require 
an  appreciable  concentration  of  OH~  for  its  precipitation.  The 
NH4+  from  the  ammonium  salts  combines  with  the  OH~  formed  by 
the  ferrous  hydroxide  to  give  molecular  ammonium  hydroxide. 
Ferrous  oxide  FeO  is  black,  and  is  formed  by  heating  ferrous  oxalate 
in  absence  of  air.  It  may  be  made  also  by  cautious  reduction  of 
ferric  oxide  by  hydrogen  (at  about  300°),  but  is  easily  reduced  further 
to  the  metal.  It  catches  fire  spontaneously  when  exposed  to  the  air. 

Ferrous  carbonate  FeCO3  is  found  in  nature  as  siderite,  and  may 
be  made  in  slightly  hydrolyzed  form  by  precipitation.  The  precipi- 
tate is  white,  but  rapidly  darkens  and  finally  becomes  brown,  the 
ferrous  hydroxide  produced  by  hydrolysis  being  oxidized  to  the  ferric 
condition.  The  salt  interacts  with  water  containing  carbonic  acid 
after  the  manner  of  calcium  carbonate  (p.  705),  giving  the  more 
soluble  Fe(HCO3)2  and  hence  the  latter  is  found  in  solution  in 
natural  (chalybeate)  waters. 

Ferrous  sulphide  FeS  may  be  formed  as  a  black,  metallic-looking 
mass  by  heating  together  the  free  elements.  It  is  produced  by  pre- 
cipitation with  ammonium  sulphide,  but  incompletely  or  not  at  all 
with  hydrogen  sulphide.  It  interacts  readily  with  dilute  acids.  The 
precipitated  form  is  slowly  oxidized  to  ferrous  sulphate  by  the  air. 

Ferrous  sulphate  FeSO4  is  obtained  by  allowing  pyrites  to  oxidize 
in  the  air  and  leaching  the  residue: 

2FeS2  +  702  +  2H20  ->  2FeS04  +  2H2S04. 

The  liquor  is  treated  with  scrap  iron  and  the  neutral  solution  evapo- 
rated until  a  hydrate  FeSO4,7H2O,  green  vitriol,  or  "  copperas,"  is 
deposited.  This  substance  forms  green  crystals  belonging  to  the 
monosymmetric  system,  but  gives  also  mixed  crystals  in  which  it  is 
isomorphous  with  the  rhombic  vitriols  (cf.  Magnesium  and  Zinc 
sulphates,  p.  771).  The  crystals  are  efflorescent,  and  also  become 
brown  from  oxidation  to  a  basic  ferric  sulphate : 

4FeS04  +  02  +  2H£0  ->  4Fe(OH)S04. 

With  excess  of  sulphuric  acid  and  an  active  oxidizing  agent,  such  as 
nitric  acid,  ferric  sulphate  is  formed.  The  double  salts  of  the  form 
(NH4)2SO4,FeSO4,6H2O  (Mohr's  salt)  are  not  efflorescent,  and  in  solid 
form  are  less  readily  oxidized  than  is  ferrous  sulphate.  The  ferroua 


IRON,   COBALT,   NICKEL  895 

sulphate  is  used  in  dyeing  and  in  making  writing  ink.  The  extract  of 
nut-galls  contains  tannic  acid  HCuHgOg,  which,  with  ferrous  sul- 
phate, gives  ferrous  tannate,  a  soluble,  almost  colorless  salt.  A 
solution  of  this  salt  containing  gum-arabic  and  some  blue  or  black 
dye  constitutes  the  ink.  When  the  writing  is  exposed  to  the  air, 
the  ferrous  tannate  is  oxidized  to  the  ferric  condition,  and  the  ferric 
compound  is  a  fine,  black  precipitate  (cf.  p.  597).  The  dye  is  added 
to  make  the  writing  visible  from  the  first.  The  black  streaks  seen 
below  nail-heads  in  oak  and  other  woods  are  due  to  the  formation  of 
ferrous  carbonate  and  its  interaction  with  the  tannic  acid  in  the  wood. 
Ferrous  sulphate  is  also  used  in  the  purification  of  water  (p.  815). 

Ferric  Compounds.  —  By  leading  chlorine  into  a  solution  of 
ferrous  chloride,  and  evaporating  until  the  proper  proportion  of  water 
alone  remains,  a  yellow,  deliquescent  hexahydrate  of  ferric  chloride, 
FeCl3,6H2O  is  obtained.  When  this  is  heated  still  further,  hydrol- 
ysis takes  place  and  the  oxide  remains.  When  chlorine  is  passed 
over  heated  iron,  anhydrous  ferric  chloride  sublimes  in  dark  green 
scales,  which  are  red  by  transmitted  light.  At  a  high  temperature  the 
formula  of  the  vapor  is  FeCl3,  but  at  lower  temperatures,  in  con- 
sequence of  association,  the  density  increases  and  the  formula  ap- 
proaches Fe2Cl6.  In  solution,  the  salt,  like  other  ferric  salts,  can  be 
reduced  to  the  ferrous  condition  by  boiling  with  iron:  2Fe~f"H~  4-  Fe 
—  »  3Fe++.  The  same  reduction  is  effected  by  hydrogen  sulphide  and 
by  stannous  chloride  (cf.  Mercuric  chloride,  p.  778)  : 

2Fe+++  +  S=  -*  2Fe++  +  S  j  . 
2Fe+++  +  Sn++  -»  2Fe++  + 


The  last  action  shows  that  ferrous  salts  are  less  active  reducing  agents 
than  are  the  stannous  salts.  The  ferric  ion  is  almost  colorless,  the 
yellow-brown  color  of  solutions  of  ferric  salts  being  due  to  the  presence 
of  ferric  hydroxide  produced  by  hydrolysis.  The  color  deepens  when 
the  solution  is  heated,  and  fades  again  very  slowly,  by  reversal  of  the 
action,  when  the  cold  solution  is  allowed  to  stand.  On  the  other 
hand,  the  hydrolysis  may  be  reversed  and  the  color  may  be  almost 
destroyed,  particularly  in  the  case  of  the  nitrate,  when  excess  of  the 
acid  is  added  to  the  solution: 

Fe(N03)3  +  3H20  ±=>  Fe(OH)3  +  3HNO3. 

Ferric  iodide  is  reduced  by  the  hydriodic  acid  produced  by  its  own 
hydrolysis,  and  hence  ferrous  iodide  does  not  unite  with  iodine  to 


896  INORGANIC   CHEMISTRY 

form  this  compound.  The  case  is  similar  to  that  of  cupric  iodide 
(p.  741). 

Ferric  hydroxide  Fe(OH)3  appears  as  a  brown  precipitate  when  a 
base  is  added  to  a  ferric  salt.  It  does  not  interact  with  excess  of  the 
alkali.  In  this  gelatinous  form  the  substance  dries  to  the  oxide  with- 
out giving  definite  intermediate  hydrated  oxides.  The  hydrates, 
Fe4O3(OH)6  (brown  iron  ore)  and  Fe20(OH)4  (bog  iron  ore),  however, 
are  found  in  nature.  The  hydroxide  passes  easily  into  colloidal  sus- 
pension in  a  solution  of  ferric  chloride,  and  by  subsequent  dialysis 
through  a  piece  of  parchment  (cf.  p.  621)  the  salt  can  be  separated, 
and  a  pure  colloidal  suspension  of  the  hydroxide  obtained.  This 
suspension,  known  as  dialyzed  iron,  is  red  in  color,  shows  no  depres- 
sion in  the  freezing-point,  and  is  not  an  electrolyte.  The  hydroxide 
is  a  positive  colloid  and  is  coagulated  (brown  precipitate)  by  the 
addition  of  salts,  bivalent  negative  ions  being  more  effective  than 
univalent  ones  (p.  844). 

Ferric  oxide  Fe2C>3  is  sold  as  "  rouge  "  and  "  Venetian  red."  It  is 
made  from  the  ferrous  sulphate  obtained  in  cleaning  ironware  which 
is  to  be  tinned  or  galvanized.  The  salt  is  allowed  to  oxidize,  and  the 
ferric  hydroxide,  thrown  down  by  the  addition  of  lime,  is  calcined. 
A  purer  form  is  produced  by  dry  distillation  of  the  basic  ferric  sul- 
phate, an  operation  which  used  to  be  undertaken  on  a  large  scale  for 
making  Nordhausen  sulphuric  acid  (p.  430).  The  product  varies 
in  tint  from  a  bright  yellowish-red  to  a  dark  violet-brown  according 
to  the  fineness  of  the  powder.  The  best  rouge  is  obtained  by  calcining 
ferrous  oxalate  FeC2O4.  This  oxide  is  not  distinctly  acidic,  but  by 
fusion  with  more  basic  oxides,  compounds  like  franklinite  Zn(Fe02)2 
may  be  formed.  It  is  reduced  by  hydrogen,  at  about  300°  to  ferrous 
oxide,  and  at  700-800°  to  metallic  iron. 

Magnetic  oxide  of  iron  Fe3O4  or  lodestone  is  found  in  nature,  and 
is  formed  by  the  action  of  air  (hammer-scale),  steam,  or  carbon 
dioxide  on  iron.  It  forms  octahedral  crystals  like  the  spinelles 
(p.  812),  and  is  a  ferrous-ferric  oxide  FeO,Fe2O3  or  Fe(FeO2)2,  related 
to  franklinite. 

Ferric  sulphide  Fe2S3  may  be  made  by  fusing  together  the  free 
elements.  It  is  obtained  by  precipitation  when  soluble  sulphides  are 
added  to  solutions  of  ferric  salts  (Stokes) : 

Fe2(SO4)3  +  3(NH4)2S  -» Fe2S3  +  3(NH4)2S04. 

Formerly  the  precipitate  was  supposed  to  be  a  mixture:  2FeS  +  S. 
Ferric  sulphate  Fe2(S04)a  is  formed  by  oxidation  of  ferrous  sul- 


IRON,  COBALT,  NICKEL  897 

phate,  and  is  obtained  as  a  white  mass  by  evaporation.  It  gives 
alums,  such  as  (NH4)2SO4,Fe2(SO4)3,24H20,  which  are  almost  color- 
less when  pure,  but  usually  have  a  pale  reddish-violet  tinge. 

Pyrite  FeS2.  —  The  mineral  pyrite  (Fools'  gold)  is  the  sulphide 
of  iron  which  is  most  stable  in  the  air.  It  is  found  in  nature  in  the 
form  of  glittering,  golden-yellow  cubes,  octahedrons,  and  pentagonal 
dodecahedrons.  It  is  not  attacked  by  dilute  acids,  but  concentrated 
hydrochloric  acid  slowly  converts  it  into  ferrous  chloride  and  sulphur. 
It  is  reduced  by  hydrogen  to  ferrous  sulphide. 

Cyanides.  —  When  potassium  cyanide  is  added  to  solutions  of 
ferrous  salts,  yellowish  precipitates  are  produced,  but  the  simple 
cyanides  cannot  be  obtained  in  pure  form.  These  precipitates  in- 
teract with  excess  of  the  cyanide,  giving  a  soluble  complex  cyanide 
of  the  form  4KCN,Fe(CN)2.  This  is  called  ferrocyanide  of  potas- 
sium. Ferric  salts  give  only  ferric  hydroxide. 

Ferrocyanide  of  potassium  K4Fe(CN)6,3H20,  "yellow  prussiate  of 
potash,"  is  made  by  heating  nitrogenous  animal  refuse,  such  as  blood, 
with  iron  filings  and  potassium  carbonate.  The  resulting  mass  con- 
tains potassium  cyanide  and  ferrous  sulphide,  and  when  it  is  treated 
with  warm  water  these  interact  and  produce  the  ferrocyanide: 

2KCN  +  FeS  ->  Fe(CN)2  -f  K2S, 
4KCN  +  Fe(CN)2  ->  K4.Fe(CN)6. 

The  salt  is  made  also  from  the  cyanogen  contained  in  crude  illumi- 
nating-gas. The  trihydrate  forms  large,  yellow,  mono&ymmetric 
tables.  The  solution  contains  almost  exclusively  the  ions  K+  and 
Fe(CN)6==,  and  gives  none  of  the  reactions  of  the  ferrous  ion  F6++. 
The  corresponding  acid  H4Fe(CN)6  may  be  obtained  as  white  crys- 
talline scales  by  addition  of  an  acid  and  of  ether  (with  which  the 
substance  forms  a  compound  or  solid  solution)  to  the  salt.  The  acid 
is  a  fairly  active  one,  but  is  unstable  and  decomposes  in  a  complex 
manner.  Other  f errocyanides  may  be  made  by  precipitation.  That  of 
copper  Cu2Fe(CN)6  is  brown,  and  ferric  ferrocyanide  Fe4[Fe(CN)6]3 
has  a  brilliant  blue  color  (Prussian  blue).  The  ferrous  compound 
(insoluble)  Fe2Fe(CN)6,  or  perhaps  K2FeFe(CN)6,  is  white,  but  quickly 
becomes  blue  by  oxidation.  The  f  errocyanides  are  not  poisonous. 

Ferricyanide  of  potassium  K3Fe(CN)6  is  easily  made  from  the 
1  ferrocyanide  by  oxidation: 

2K4Fe(CN)6  +  C12  ->  2KC1  +  2K3Fe(CN)6, 
or  2Fe(CN)6==  +  C12  -*  2Fe(CN)6=-  +  2C1-. 


898  INORGANIC   CHEMISTRY 

It  forms  red  monosymmetric  prisms.  The  free  acid  H3Fe(CN)6  is 
unstable.  Other  salts  may  be  prepared  by  precipitation.  Ferrous 
ferricyanide  Fe3[Fe(CN)6]2  is  deep-blue  in  color  (TurnbulTs  blue). 
With  ferric  salts  only  a  brown  solution  is  obtained. 

Prussian  blue  and  TurnbulPs  blue  are  used  in  making  laundry 
blueing.  They  are  insoluble,  but  give  colloidal  suspensions  and  are 
adsorbed  by  the  material  of  the  cloth. 

Ferric  thiocyanate  Fe(NCS)3  is  formed  by  interaction  of  soluble 
thiocyanates  with  ferric  salts  (cf.  p.  292).  It  is  deep-red  in  color  and 
gives  a  blood-red  solution  in  water.  Since  both  the  ions  are  colorless, 
the  solution  must  contain  much  of  the  molecular  salt.  Its  formation 
furnishes  a  very  delicate  test  for  traces  of  ferric  salts. 

Blue-Prints.  —  Some  ferric  salts,  when  exposed  to  light,  are 
reduced  to  the  ferrous  condition.  Thus,  ferric  oxalate,  in  the  light, 
gives  ferrous  oxalate: 

Fe2(C2O4)3  ->  2FeC2O4  +  2CO2. 

When  paper  is  coated  with  ferric  oxalate  solution  and  dried,  and 
an  ink  drawing  on  transparent  paper  is  placed  over  the  prepared 
surface,  sunlight  will  reduce  the  iron  to  the  ferrous  condition,  Except- 
ing where  the  ink  protects  it.  When  the  sheet  is  then  dipped  in 
potassium  ferricyanide  solution  (developer),  the  ferric  oxalate  gives 
only  the  brown  substance  which  can  be  washed  out.  But  the  deep 
blue,  insoluble  ferrous  ferricyanide  is  precipitated  in  the  pores  of  the 
paper  where  the  light  has  acted.  The  drawing  appears  white  on  a 
blue  blackground.  In  ordinary  blue-print  paper,  ammonium-ferric 
citrate  takes  the  place  of  the  oxalate,  and  the  ferricyanide  has  already 
been  applied  to  the  paper  before  drying,  so  that  only  exposure  and 
washing  remain  to  be  done.  Dilute  sodium  hydroxide  solution 
decomposes  the  ferricyanide,  and  is  used  for  writing  (in  white)  on 
blue-prints. 

Iron  Carbonyls.  —  When  carbon  monoxide  is  led  over  finely 
divided  iron  at  40-80°,  or  under  eight  atmospheres  pressure  at  the 
ordinary  temperature,  volatile  compounds  of  the  composition  Fe(CO)4 
iron  tetracarbonyl,  and  Fe(CO)5,  the  pentacarbonyl,  are  formed. 
When  the  gaseous  mixture  is  heated  more  strongly,  the  compounds 
decompose  again,  and  iron  is  deposited.  Illuminating-gas  burners 
frequently  receive  a  deposit  of  iron  from  this  cause. 


IRON,   COBALT,   NICKEL  899 

Ferrates.  —  A  red  solution  of  potassium  ferrate  K2FeO4,  is  ob- 
tained by  passing  chlorine  through  caustic  potash  in  which  ferric 
hydroxide  is  suspended,  or  by  heating  pulverized  iron  with  dry  sodium 
nitrate.  The  salt  crystallizes  in  red,  rhombic  prisms,  isomorphous 
with  the  sulphate  and  chromate  of  potassium.  It  alters  quickly  in 
solution,  in  consequence  of  hydrolysis  and  subsequent  decomposition 
of  the  ferric  acid,  depositing  ferric  hydroxide  and  giving  off  oxygen. 
Barium,  strontium,  and  calcium  salts  are  formed  as  red  precipitates 
by  double  decomposition. 

Analytical  Reactions  of  Compounds  of  Iron.  —  There  are 
two  ionic  forms  of  iron,  ferrous-ion  Fe"1"1",  which  is  very  pale-green, 
and  ferric-ion  Fe+++,  which  is  almost  colorless.  The  yellow  color  of 
ferric  salts  is  due  to  hydrolysis.  Ammonium  sulphide  gives  with  the 
former  black  ferrous  sulphide,  which  is  soluble  in  dilute  acids.  The 
hydroxides  are  white  and  brown  respectively,  and  ferrous  carbonate 
is  white.  With  ferric  salts,  which  are  hydrolyzed  (about  5  per  cent), 
soluble  carbonates  yield  the  hydroxide,  because  they  neutralize  the 
free  acid  and  displace  the  equilibrium.  With  ferrocyanide  of  potas- 
sium, ferrous  salts  give  a  white,  and  ferric  salts  a  blue,  precipitate. 
With  ferricyanide  of  potassium  the  former  give  a  deep-blue  precipi- 
tate, and  the  latter  a  brown  solution.  Ferric  thiocyanate  is  deep-red. 
From  ferric  solutions  barium  carbonate  throws  down  ferric  hydroxide. 
When  sodium  acetate  is  added  in  excess  to  a  ferric  salt,  a  red,  little 
ionized,  but  easily  hydrolyzed,  ferric  acetate  is  formed.  When  the 
solution  is  boiled  the  hydrolysis  is  increased,  and  an  insoluble,  basic 
ferric  acetate  is  thrown  down.  With  borax,  iron  compounds  give  a 
bead  which  is  green  (ferrous  borate)  in  the  reducing  flame,  and 
colorless  or,  with  much  iron,  yellow  (ferric  borate)  or  even  brown 
when  oxidized.  *P. 

COBALT  Co 

The  Chemical  Relations  of  the  Element.  —  Cobalt  forms 
cobaltous  and  cobaltic  oxides  and  hydroxides  CoO  and  Co(OH)2, 
CosOs  and  Co(OH)3/  respectively,  which  are  all  basic,  the  former 
more  so  than  the  latter.  The  cobaltous  salts  are  little  hydrolyzed, 
but  the  cobaltic  salts  are  largely  decomposed  by  water.  The  latter 
also  liberate  readily  one-third  of  the  negative  radical,  after  the  man- 
ner of  manganic  salts,  becoming  cobaltous.  Complex  cations  and 
anions  containing  cobalt  are  very  numerous  and  very  stable. 


900  INORGANIC  CHEMISTRY 

Occurrence  and  Properties.  —  Cobalt  is  found  along  with 
nickel  in  smaltite  CoAs2  and  cobaltite  CoAsS.  The  pure  metal  may 
be  made  by  Goldschmidt's  process,  or  by  reducing  the  oxalate,  or  an 
oxide,  with  hydrogen. 

The  metal  is  silver-white,  with  a  faint  suggestion  of  pink.  It  is 
less  tough  than  iron,  and  has  no  commercial  applications.  It  dis- 
places hydrogen  slowly  from  dilute  acids,  but  interacts  readily  with 
nitric  acid. 

Cobaltous  Compounds.  —  The  chloride  CoCl2,6H2O  may  be 
made  by  treating  the  oxide  with  hydrochloric  acid.  It  forms  red 
prisms,  and  when  partially  or  completely  dehydrated  becomes  deep- 
blue.  Writing  made  with  a  diluted  solution  upon  paper  is  almost 
invisible,  but  becomes  blue  when  warmed  and  afterwards  takes  up 
moisture  from  the  air,  and  is  once  more  invisible  (sympathetic 
ink).  Most  cobaltous  compounds  are  red  when  hydrated  or  in  solu- 
tion (Co++)  and  blue  when  dehydrated.  The  blue  color  assumed  by  a 
strong  solution  of  cobaltous  chloride,  when  it  is  warmed,  or  when 
hydrochloric  acid  is  added  to  it,  is  explained  by  some  chemists  as 
being  due  to  repression  of  the  ionization  of  the  salt,  and  by  others 
as  being  due  to  the  formation  of  the  complex  anion  of  the  salt 
Co++.CoCl4~.  By  addition  of  sodium  hydroxide  to  a  cobaltous  salt, 
a  blue  basic  salt  is  precipitated.  When  the  mixture  is  boiled,  the 
pink  cobaltous  hydroxide,  Co(OH)2  is  formed.  This  becomes  brown 
through  oxidation  by  the  air.  It  interacts  with  ammonium  hydrox- 
ide, giving  a  soluble  ammonio-cobaltous  hydroxide  (cf.  p.  784),  which 
is  quickly  oxidized  by  the  air  to  an  ammonio-cobaltic  compound  (see 
below) .  It  interacts  also  with  salts  of  ammonium  as  does  magnesium 
hydroxide  (p.  766).  When  dehydrated  it  leaves  the  black  cobaltous 
oxide.  Cobaltous  sulphate  CoSO4,7H2O,  isomorphous  with  mag- 
nesium sulphate,  and  the  nitrate  Co(NO3)2,6H20  are  familiar  salts. 
The  black  cobaltous  sulphide  CoS  is  precipitated  by  ammonium 
sulphide  from  solutions  of  all  salts,  and  even  by  hydrogen  sulphide 
from  the  acetate,  or  a  solution  containing  much  sodium  acetate  (cf. 
p.  716).  Once  it  has  been  formed,  it  does  not  interact  even  with 
dilute  hydrochloric  acid,  having  apparently  changed  into  a  less 
active  form.  A  sort  of  cobalt  glass,  made  by  fusing  sand,  cobalt 
oxide,  and  potassium  nitrate,  forms,  when  powdered,  a  blue  pigment, 
smalt,  used  in  china-painting  and  by  artists. 

Cobaltic  Compounds.  —  By  addition  of  a  bypochlorite  to  a 
solution  of  a  cobaltous  salt,  cobaltic  hydroxide  Co(OH)3,  a  black 


IRON,   COBALT,  NICKEL  901 

powder,  is  precipitated.  Cautious  ignition  of  the  nitrate  gives 
cobaltic  oxide  Co2O3.  Stronger  ignition  gives  the  commercial  oxide 
which  is  a  cobalto-cobaltic  oxide  Co3O4.  Cobaltic  oxide  dissolves 
in  cold  hydrochloric  acid,  but  the  solution  gives  off  chlorine  when 
warmed.  By  placing  cobaltous  sulphate  solution  round  the  anode 
of  an  electrolytic  cell,  crystals  of  cobaltic  sulphate  CosCSO^s  have 
been  made  and  cobaltic  alums  have  also  been  prepared  (Hugh 
Marshall). 

Complex  Compounds.  —  Potassium  cyanide  precipitates  from 
cobaltous  salts  a  brownish-white  cyanide.  This  interacts  with  ex- 
cess of  the  reagent,  giving  a  solution  of  potassium  cobaltocyanide 
K4Co(CN)e  (cf.  p.  897).  This  compound  is  easily  oxidized  by  chlo- 
rine, or  even  when  the  solution  is  boiled  in  the  air,  and  the  colorless 
potassium  cobalticyanide  is  formed: 

4K4Co(CN)6  +  2H20  +  02  ->  4K3.Co(CN)6  +  4KOH. 

The  solution  gives  none  of  the  reactions  of  Co~l~H~,  and  with  acids  the 
very  stable  cobalti cyanic  acid  H3Co(CN)3  is  liberated. 

When  acetic  acid  and  potassium  nitrite  are  added  to  a  cobaltous 
salt  the  latter  is  oxidized  by  the  nitrous  acid  (liberated  by  the  acetic 
acid)  and  a  yellow  complex  salt  K3.Co(NO2)6,nH2O  (=  Co(N02)3,- 
3KNO2),  potassium  cobaltinitrite,  is  thrown  down. 

Cobaltic  salts  give  with  ammonia  complex  compounds  which  are 
many  and  various.  The  cations  often  contain  negative  groups,  and 
are  such  as  Co(NH3)6+++,  Co(NH3)5Cl++  and  Co(NH3)6NO2++.  Usu- 
ally the  solutions  give  none  of  the  reactions  of  cobaltic  ions,  and  often 
fail  likewise  to  give  those  of  the  anion  of  the  original  salt. 

NICKEL  Ni 

The  Chemical  Relations  of  the  Element.  —  Nickel  forms 
nickelous  and  nickelic  oxides  and  hydroxides  NiO  and  Ni(OH)2,  Ni2O3 
and  Ni(OH)3,  but  only  the  former  are  basic.  The  nickelous  salts 
resemble  the  cobaltous  and  ferrous  salts,  but  are  not  oxidizable 
into  corresponding  nickelic  compounds.  Since  there  are  no  nickelic 
salts,  there  are  here  no  analogues  of  the  cobalticyanides  or  the  cobal- 
tinitrites.  The  complex  nickelous  salts,  like  the  complex  cobaltous 
salts,  and  unlike  the  complex  cobaltic  salts,  are  unstable,  and  so  give 
some  of  the  reactions  of  Ni++ 


902  INORGANIC  CHEMISTRY 

Occurrence  and  Properties. — Nickel  occurs  in  meteorites  (free) 
and  in  niccolite  NiAs  and  nickel  glance  NiAsS.  It  is  now  manu- 
factured chiefly  from  pentlandite  [Ni,Cu,Fe]S  and  other  minerals 
found  at  Sudbury,  Ontario  (production  in  1913,  24,840  short  tons), 
and  from  garnierite,  a  silicate  of  nickel  and  magnesium,  found  in 
New  Caledonia.  In  the  former  case,  the  ore  is  roasted,  smelted,  and 
finally  bessemerized.  The  resulting  alloy  of  copper  and  nickel  is 
much  used  for  sheet-metal  work  (Monel  metal,  approx.  1  : 1),  Pure 
nickel  is  separated  from  the  copper  by  an  electrolytic  process  (p.  747), 
or  by  the  Monde  process  (see  below). 

The  metal  is  white,  with  a  faint  tinge  of  yellow,  is  very  hard;  and 
takes  a  high  polish  (m.-p.  1452°).  It  is  used  in  making  alloys,  such 
as  German  silver  (copper,  zinc,  nickel,  2:1:1)  and  the  "nickel" 
used  in  coinage  (copper,  nickel,  3  : 1).  Although  in  these  alloys 
the  red  color  of  the  copper  is  completely  lost,  the  copper  is  simply 
dissolved,  and  not  combined.  Zinc  and  copper,  however,  give  a 
compound  Cu2Zn3.  Nickel  plating  on  iron  is  accomplished  exactly 
like  silver  plating  (p.  755).  The  bath  contains  an  ammoniacal  solu- 
tion of  ammonium-nickel  sulphate  (NH^SO^NiSO^GH^O,  and  a 
plate  of  nickel  forms  the  anode. 

The  metal  rusts  very  slowly  in  moist  air.  It  displaces  hydrogen 
with  difficulty  from  dilute  acids,  but  interacts  with  nitric  acid. 

Compounds  of  Nickel.  —  The  chloride  NiCl2,6H2O  is  made  by 
treating  one  of  the  oxides  with  hydrochloric  acid,  and  is  green  in 
color  (when  anhydrous,  brown).  The  sulphate  NiSO4,6H20,  which 
crystallizes  in  green,  square  prismatic  forms  at  30-40°,  is  the  most 
familiar  salt.  The  heptahydrate  NiSO4,7H2O,  obtained  from  cold 
solutions,  is  isomorphous  with  magnesium  sulphate.  Nickelous  hy- 
droxide Ni(OH)2  is  formed  as  an  apple-green  precipitate,  and  when 
heated  leaves  the  green  nickelous  oxide  NiO.  It  interacts  with 
ammonium  hydroxide,  giving  a  complex  ammonio-nickel  cation.  It 
also  interacts  with  and  dissolves  in  salts  of  ammonium  (cf.  p.  766). 
By  cautious  ignition  of  the  nitrate,  nickelic  oxide  Ni203,  is  formed  as  a 
black  powder.  The  oxides  and  salts,  when  heated  strongly  in  oxygen, 
give  the  oxide  Ni3O4.  The  last  two  oxides  liberate  chlorine  when 
treated  with  hydrochloric  acid,  and  give  nickelous  chloride.  Nickelic 
hydroxide  Ni(OH)3  is  a  black  precipitate  formed  when  a  hypochlorite 
is  added  to  any  salt  of  nickel.  Nickelous  sulphide  is  thrown  down  by 
ammonium  sulphide,  and  behaves  like  cobaltous  sulphide  (p.  900). 


IRON,   COBALT,  NICKEL  903 

It  forms  a  brown  colloidal  solution  when  excess  of  the  precipitant  is 
usedj  and  is  then  deposited  very  slowly. 

Addition  of  dimethylglyoxime  to  an  ammoniacal  solution  of  a 
salt  of  nickel  gives  a  scarlet  precipitate  of  an  acid  salt: 

Ni(OH)2  +  2(HON)2C2(CH3)2  -> 2H20  +  NiH2[(ON)2C2(CH3)2]2. 
This  reaction  is  not  shown  by  salts  of  cobalt,  especially  if  oxidation 
to  the  cobaltic  condition  has  been  permitted  by  contact  with  air. 

With  potassium  cyanide  and  a  salt  of  nickel  the  greenish  nickelous 
cyanide  Ni(CN)2  is  first  precipitated.  This  dissolves  in  excess  of  the 
reagent,  and  a  complex  salt  K2Ni(CN)4,H2O  (=  2KCN,Ni(CN)2) 
may  be  obtained  from  the  solution.  This  salt  is  of  different  composi- 
tion from  the  corresponding  compounds  of  cobalt  and  iron,  and  is  less 
stable.  Thus,  with  bleaching  powder,  it  gives  Ni(OH)3  as  a  black 
precipitate.  When  the  solution  is  boiled  in  the  air  no  oxidation  to  a 
complex  nickelicyanide  occurs,  and  indeed  no  such  salts  are  known. 
This  fact  enables  the  chemist  to  separate  cobalt  and  nickel,  for  when 
the  mixed  cyanides  are  boiled  and  then  treated  with  bleaching  pow- 
der, the  cobalticyanide  is  unaffected.  With  potassium  nitrite  and 
acetic  acid  no  insoluble  compound  corresponding  to  that  given  by 
cobalt  salts  is  formed  by  salts  of  nickel.  The  only  known  compound 
which  could  be  formed,  4KNO2,Ni(NO2)2,  is  soluble.  This  action 
also  is  used  for  the  purpose  of  separation.  The  pink  color  of  cobalt 
salts  and  the  green  of  nickel  salts  are  complementary  colors,  so  that, 
by  using  suitable  proportions  of  the  two,  a  colorless  mixture  can  be 
produced. 

When  finely  divided  nickel,  made  by  reducing  the  oxide  or  oxalate 
with  hydrogen  at  a  moderate  temperature,  is  exposed  to  a  stream  of 
cold  carbon  monoxide,  nickel  carbonyl  Ni(CO)4  is  formed.  This  is  a 
vapor  and  is  condensable  to  a  colorless  liquid  (b.-p.  43°  and  m.-p. 

—  25°).     The  vapor  is  poisonous.     When  heated  to  150-180°  it  is  dis- 
sociated and  nickel  is  deposited.     Cobalt  forms  no  corresponding 
compound.      In  commerce,  pure  nickel  is  separated  from   copper 
(and  cobalt)  in  the  Monde  process  by  passing  carbon  monoxide  over 
the  pulverized  alloy,  and  subsequently  heating  the  gas. 

Analytical  Reactions  of  Compounds  of  Cobalt  and  Nickel. 

—  Cobaltous-ion  Co++  is  pink,  and  the  nickelous  ion  Ni++  green. 
The  reactions  used  in  analysis  have  been  described  in  the  preceding 
paragraphs.     With  borax,  cobalt  compounds  give  a  blue  bead  (cobal- 
tous  borate),  and  nickel  compounds  a  bead  which  is  brown  in  the  oxi- 


904  INORGANIC   CHEMISTRY 

dizing  flame  and  cloudy,  from  the  presence  of  metallic  nickel,  when 
reduced. 

Exercises,  —  1.   What  would  be  the  interactions  of  calcium  car- 
bonate when  fused  with  sand  and  with  clay,  respectively? 

2.  Make  equations  representing,    (a)   the  oxidation  of  ferrous 
chloride  by  air,  (6)  the  hydrolysis  of  ferrous  carbonate  and  the  oxida- 
tion of  ferrous  hydroxide,  (c)  the  oxidation  of  ferrous  sulphate  with 
excess  of  sulphuric  acid  by  hypochlorous  acid,  (d)  the  formation  of 
ferrous  and  ferric  tannates  (p.  895),  (e)  the  reduction  of  ferric  chloride 
by  iron,  by  hydrogen  sulphide,  and  by  stannous  chloride,  respectively, 
(/)  the  dry  distillation  of  basic  ferric  sulphate,  (g)  the  formation  of 
ferric  ferrocyanide  and  of  ferrous  ferricyanide,  (h)  the  hydrolysis  and 
decomposition  of  potassium  ferrate. 

3.  Explain  the  solubility  of  cobaltous  and  nickelous  hydroxides 
in  salts  of  ammonium. 

4.  Construct  equations  to  show  the  formation,  (a)  of  the  insoluble 
potassium  cobaltini trite  (nitric  oxide  is  given  off),   (6)  of  nickelic 
hydroxide  from  nickelous  chloride  and  sodium  hypochlorite.     Re- 
membering that  the  hypochlorite  is  somewhat  hydrolyzed,  explain 
why  the  precipitation  in  (b)  is  complete. 

5.  Tabulate  in  detail  the  chemical  relations  of  the  elements  cobalt 
and  nickel,  especially  to  show  the  resemblances  and  differences. 


CHAPTER  XLV 
THE   PLATINUM  METALS 

THE  remaining  elements  of  Mendelejeff's  eighth  group  divide 
themselves  into  two  sets  of  three  each.  Just  as  iron,  cobalt,  and 
nickel  have  similar  atomic  weights  and  much  the  same  density 
(7.8-8.8),  so  ruthenium  (Ru,  at.  wt.  101.7),  rhodium  (Rh,  at.  wt.  103), 
and  palladium  (Pd,  at.  wt.  106.7)  have  densities  from  12.26  to  11.5. 
Similarly  osmium  (Os,  at.  wt.  191),  iridium  (Ir,  at.  wt.  193),  and 
platinum  (Pt,  at.  wt.  195.2)  form  a  triad  with  specific  gravities  from 
22.5  to  21.5.  Chemically,  ruthenium  shows  the  closest  resem- 
blance to  osmium,  and  both  are  allied  to  iron.  Similarly,  rhodium 
and  iridium,  and  palladium  and  platinum  are  natural  pairs. 

The  six  elements  are  found  alloyed  in  nuggets  and  particles  which 
are  separated  from  alluvial  sand  by  washing.  Platinum  forms  60-84 
per  cent  of  the  whole.  The  chief  deposits  are  in  the  Ural  Mountains, 
smaller  amounts  being  found  in  California,  Australia,  Borneo,  and 
elsewhere.  The  components  are  separated  by  a  complex  series  of 
chemical  operations. 

Ruthenium  and  Osmium.  —  These  metals  are  gray  like  iron, 
while  the  other  four  are  whiter  and  more  like  cobalt  and  nickel.  They 
also  resemble  iron  in  being  the  most  infusible  members  of  their 
respective  sets.  Both  melt  considerably  above  2000°.  They  like- 
wise resemble  iron  in  uniting  easily  with  free  oxygen,  while  the  other 
four  elements  do  not.  Ruthenium  gives  Ru2O  in  air,  and  Ru02  in 
oxygen.  Osmium  gives  OsO4,  "  osmic  acid,"  a  white  crystalline 
body  melting  at  40°  and  boiling  at  about  100°.  The  odor  and  irri- 
tating effects  of  the  vapor  recall  chlorine  (Gk.,  <xr/«},  odor).  The 
substance  is  not  an  acid,  nor  even  an  acid  anhydride.  The  aqueous 
solution  is  used  in  histology,  and  stains  tissues  in  consequence  of  its 
reduction  by  organic  bodies  to  metallic  osmium.  It  is  affected  par- 
ticularly by  fat.  Osmic  acid  also  hardens  the  material  without  dis- 
torting it.  Osmium  forms  also  a  yellow,  crystalline  fluoride  OsF8 
(m.-p.  34.5°).  Ruthenium  and  osmium  have  a  maximum  valence 
of  eight. 

905 


906  INORGANIC   CHEMISTRY 

These  elements  resemble  iron  in  giving  ruthenates  and  osmates, 
like  K2Ru04  and  K2Os04,  but  no  corresponding  oxide  or  acid. 
Potassium  ruthenate  resembles  potassium  manganate  and  gives, 
when  diluted,  an  oxide  of  ruthenium  and  potassium  perruthenate 
KRu04.  There  are  salts  derived  from  the  oxides  Ru203  and 


Rhodium  and  Iridium.  —  These  metals  are  not  attacked  by 
aqua  regia,  while  the  other  four  are  dissolved,  more  or  less  slowly. 
They  are  harder  than  platinum,  and  iridium  is  alloyed  with  this 
metal  for  the  purpose  of  increasing  its  hardness  (pens),  and  its  re- 
sistance to  the  action  of  fluorine.  They  resemble  cobalt  in  having 
no  acid-forming  properties.  Salts  corresponding  to  the  oxides  Rh203 
and  Ir2O3  are  formed,  and  iridium  gives  compounds  derived  from 
Ir02  as  well.  The  most  familiar  compounds  of  iridium  are  the  com- 
plex chlorides  X3IrCl6  (=  3XCl,IrCl3)  and  X2IrCl6  (=  2XCl,IrCl4). 
The  solutions  of  the  latter  are  red,  and  the  acid,  chloro-iridic  acid 
HJrCle,  is  often  found  in  commercial  chloroplatinic  acid  H2PtCl6, 
and  confers  upon  it  a  deeper  color. 

Palladium  and  Platinum.  —  Palladium  is  the  only  metal  of 
this  family  which  is  attacked  by  nitric  acid.  Palladium  and  platinum 
forms  -ous  and  -4c  compounds  of  the  forms  PdX2  and  PdX4,  respec- 
tively. The  oxides  PdO  and  PtO  and  the  corresponding  hydroxides 
are  basic.  When  quadrivalent,  the  metals  appear  chiefly  in  complex 
compounds,  like  H2.PtCl6,  H2.PdCl6,  in  which  the  metal  is  in  the 
anion.  Platinum  gives  also  platinates  derived  from  the  oxide  PtO2, 
and  quadrivalent  platinum  furnishes  no  well  defined  salts  in  which  it 
constitutes  by  itself  the  positive  ion. 

Palladium.  —  This  metal  (m.-p.  1549°),  named  from  the  planet- 
oid Pallas,  is  noted  chiefly  for  its  great  tendency  to  absorb  hydrogen. 
When  finely  divided,  it  takes  up  about  800  times  its  own  volume. 
The  amount  absorbed  varies  continuously  with  the  concentration 
(pressure)  of  the  hydrogen,  although  not  according  to  a  uniform 
rule,  and  the  product  is  in  part  at  least  a  solid  solution.  When  a 
strip  of  palladium  is  made  the  cathode  of  an  electrolytic  cell,  over 
900  volumes  of  hydrogen  may  be  occluded.  This  absorbed  hydrogen, 
in  consequence  of  the  catalytic  influence  of  the  metal,  reacts  more 
rapidly  than  does  the  gas,  and  consequently  a  strip  of  hydrogenized 
palladium  will  quickly  precipitate  copper  and  other  metals  less  electro- 
positive than  hydrogen  and  will  reduce  ferric  and  other  reducible  salts  : 


THE  PLATINUM   METALS  907 

CuS04  +  H2  ->  H2SO4  +  Cu,        or      Cu++  +  H2  ->  2H+  +  Cu. 

or   2Fe+++  +  H2  -»  2Fe++  +  2H+. 


The  palladious  salts  are  such  as  PdCl2,  PdSO4,  Pd(NO3)2.  Palla- 
dic  chloride,  formed  by  treating  the  metal  with  aqua  regia,  is  con- 
tained in  the  solution  in  the  form  of  chloropalladic  acid  H^PdCU 
(=  2HCl,PdCl4)  and  gives  difficultly  soluble  salts  like  K2PdCl6. 
When  the  solution  of  the  acid  is  boiled,  however,  chlorine  is  given  off 
and  PdCl2  or  H2PdCLi  remains  in  solution. 

Platinum.  —  This  metal  (dim.  of  Sp.  plata,  silver)  is  grayish- 
white  in  color,  and  is  very  ductile.  At  a  red  heat  it  can  be  welded.  It 
does  not  melt  in  the  Bunsen  flame,  but  fuses  easily  in  the  oxyhydrogen 
jet  (m.-p.  1755°).  On  account  of  its  very  small  chemical  activity 
it  is  used  in  electrical  apparatus  and  for  making  wire,  foil,  crucibles 
and  other  vessels  for  use  in  laboratories.  It  unites,  however,  with 
carbon,  phosphorus,  and  silicon,  becoming  brittle,  and  forms  fusible 
alloys  with  metals  like  antimony  and  lead.  Hence  care  has  to  be 
taken  not  to  heat  in  vessels  made  of  it  compounds  from  which  these 
elements  may  be  liberated.  It  also  interacts  with  fused  alkalies, 
giving  platinates,  but  the  alkali  carbonates  may  be  melted  in  vessels 
of  platinum  with  impunity.  The  oxygen  acids  are  without  action 
upon  it,  but  on  account  of  the  tendency  to  form  the  extremely  stable 
complex  ion  PtCl6=  (p.  537),  the  free  chlorine  and  chloride-ion  in 
aqua  regia  convert  it  into  chloroplatinic  acid  H2PtCl6. 

The  metal  condenses  oxygen  upon  its  surface  and  it  dissolves 
hydrogen.  The  finely  divided  forms  of  the  metal,  such  as  platinum 
sponge  made  by  igniting  ammonium  chloroplatinate  (NH4)2PtCle, 
and  platinum  black  made  by  adding  zinc  to  chloroplatinic  acid,  show 
this  behavior  very  conspicuously.  They  cause  instant  explosion  of 
a  mixture  of  oxygen  and  hydrogen,  in  consequence  of  the  heat  devel- 
oped by  the  rapid  union  of  that  part  of  the  gases  which  is  condensed 
in  the  metal.  A  heated  spiral  of  fine  platinum  wire  will  continue  to 
glow  if  immersed  in  the  mixture  of  alcohol  vapor  and  air  (oxygen) 
formed  by  placing  a  little  alcohol  in  a  beaker.  Some  cigar-lighters 
work  on  this  principle.  The  heat  is  developed  by  the  interaction 
between  the  substances,  which  takes  place  with  great  speed  at  the 
surface  of  the  platinum  (cf.  p.  596). 

Platinum  is  the  only  otherwise  suitable  single  substance  which  has 
the  same  coefficient  of  expansion  as  glass,  and  it  was  consequently 
fused  into  incandescent  bulbs  and  furnished  the  electrical  connection 


908  INORGANIC  CHEMISTRY 

with  the  filament  in  the  interior.  Recently  it  has  been  displaced  by 
Eldred's  wire,  containing  a  core  of  nickel  steel  with  a  jacket  of  copper 
and  an  outer  sheath  of  platinum.  This  shrinks  slightly  less  than 
does  glass  on  cooling.  Large  amounts  are  also  consumed  in  photog- 
raphy and  by  dentists.  It  is  used  in  making  jewelry,  and  in  Russia 
for  coinage.  The  price  of  the  metal  is  very  variable,  since  a  rainy 
season  in  the  Caucasus  will  render  larger  amounts  accessible  to  the 
miners,  but,  on  the  whole,  the  many  applications  which  have  been 
found  for  it  have  quintupled  its  price  in  the  last  twenty  years. 

Compounds  of  Platinum.  —  Platinous  chloride  PtCl2  is  made 
by  passing  chlorine  over  finely  divided  platinum  at  240-250°  or  by 
heating  chloroplatinic  acid  to  the  same  temperature.  It  is  greenish 
and  insoluble  in  water,  but  forms  with  hydrochloric  acid  the  soluble 
chloroplatinous  acid  H2PtCl4.  Potassium  chloroplatinite  K2PtCl4  is 
used  in  making  platinum  prints  (cf.  p.  757).  Bases  precipitate  black 
platinous  hydroxide  Pt(OH)2  which  interacts  with  acids  but  not  with 
bases.  Gentle  heating  gives  the  oxide  PtO  and  stronger  heating 
the  metal.  With  potassium  cyanide  and  barium  cyanide  soluble 
platino-cyanides,  K2Pt(CN)4,3H2O  and  BaPt(CN)4,4H20,  are  formed. 
These  substances,  when  solid,  show  strong  fluorescence,  converting 
X-rays  as  well  as  ultra-violet  rays  into  visible  radiations.  The  barium 
salt  is  used  on  screens  to  receive  the  shadows  cast  by  X-rays. 

Platinic  chloride  PtCl4  may  be  made  by  treating  the  metal  with 
aqua  regia  and  heating  the  chloroplatinic  acid  H2PtCle  so  formed  in 
a  stream  of  chlorine  at  360°.  When  dissolved  in  water,  however,  it 
gives  a  compound  H2.PtCl4O  in  which  platinum  is  in  the  anion.  Its 
solution  deposits  red,  non-deliquescent  crystals  of  H2PtCl4O,4H2O. 

Chloroplatinic  acid  forms  reddish-brown  deliquescent  crystals 
H2PtCl6,6H2O.  With  potassium  and  ammonium  salts,  it  yields  the 
sparingly  soluble,  yellow  chloroplatinates  K2PtCl6  and  (MH4)2PtCl6 
(cf.  p.  677),  in  solutions  of  which  the  platinum  migrates  towards  the 
anode  and  silver  salts  precipitate  Ag2PtCl6  and  not  silver  chloride. 

Bases  interact  with  chloroplatinic  acid,  giving  a  yellow  or  brown 
precipitate  of  platinic  hydroxide  Pt(OH)4.  This  substance  interacts 
both  with  acids  and  with  bases.  In  the  latter  case  platinates,  like 
Na2HioPt3Oi2,H2O,  have  been  obtained.  Both  sets  of  platinum  com- 
pounds interact  with  hydrogen  sulphide,  giving  the  sulphides,  PtS 
and  PtS2.  These  are  black  powders  which  interact  with  yellow 
ammonium  sulphide  solution  giving  ammonium  sulphoplatinates. 


APPENDIX 

I.   The  Metric  System 

Length.  1  meter  (1  m.)  =  10  decimeters  =  100  centimeters  (100 
cm.)  =  1000  millimeters  (1000  mm.). 

1  kilometer  =  1000  meters  (1000  m.). 

1  decimeter  =  0.1  m.  =  10  centimeters  =  3.937  inches. 

1  meter  =  1.094  yards  =  3.286  ft.  =  39.37  inches. 

Volume.  1  liter  =  1000  cubic  centimeters  (1000  c.c.)  =  a 
cube  10  cm.  X  10  cm.  X  10  cm. 

1  liter  =  0.03532  cu.  ft.  =  61.03  cu.  in.  =  1.057  quarts  (U.  S.)  or 
1.136  quarts  (Brit.)  =  34.1  fl.  oz.  (U.  S.)  =  35.3  oz.  (Brit.). 

1  fl.  ounce  (U.  S.)  =  29.57  c.c.     1  ounce  (Brit.)  =  28.4  c.c. 
1  cu.  ft.  =  28.32  liters. 

Weight.  1  gram  (g.)  =  wt.  of  1  c.c.  of  water  at  4°  C.  1  kilo- 
gram =  1000  g. 

1  gram  =  10  decigrams  =  100  centigrams  (100  cgm.)  =  1000 
milligrams  (1000  mgm.). 

1  kilogram  =  2.205  Ibs.  avoird.  (U.  S.  and  Brit.). 

1  Ib.  avoird.  =  453.6  g. 

1  oz.  avoird.  (U.  S.  and  Brit.)  =  28.35  g.     100  g.  =  3.5  oz. 

1  nickel  (U.  S.)  weighs  5  g,     1  halfpenny  (Brit.),  5  to  5.7  g. 

1  long  ton  =  2240  Ibs.  1  short  ton  =  2000  Ibs.  1  metric  ton  = 
1000  kilos  =  2205  Ibs. 

II.   Scale  of  Hardness 

Each  of  the  following  minerals  will  scratch  the  surface  of  a  speci- 
men of  any  one  preceding  it  in  the  list. 

1.  Talc  6.  Felspar 

2.  Gypsum  (or  NaCl)  7.  Quartz 

3.  Calcite  (or  Cu)  8.  Topaz 

4.  Fluorite  9.  Corundum 

5.  Apatite  10.  Diamond 

Glass  is  slightly  scratched  by  5,  and  easily  by  those  following. 
Glass  will  not  scratch  5  distinctly,  but  will  scratch  those  preceding  5. 
A  good  knife  scratches  6  slightly,  but  not  those  following. 
A  file  will  scratch  7,  but  not  those  following. 

909 


910 


APPENDIX 


III.   Temperatures  Centigrade  and  Fahrenheit 

Upon  the  centigrade  scale,  the  freezing-point  of  water  is  0°  C. 
and  the  boiling-point  100°  C.  Upon  the  Fahrenheit  scale,  the 
same  points  are  32°  F.  and  212°  F.,  respectively.  The  same  inter- 
val is  thus  100°  on  the  one  scale  and  180°  on  the  other.  The  degree 
Fahrenheit  is  therefore  il#  or  f  of  1°  Centigrade.  Any  tempera- 
tures can  be  converted  by  using  the  formulae: 

C.°  =  I  (F.°  -  32),  F.°  =  t  (C.°)  +  32. 

The  following  table  (IV)  contains  the  temperatures  from  0°  C. 
to  35°  C.,  with  the  corresponding  values  on  the  Fahrenheit  scale 
(32°  F.  to  95°  F.). 

IV.   Vapor  Pressures  of  Water 

Both  the  Fahrenheit  (F.)  or  ordinary  and  the  Centigrade  (C.)  temperatures  are  given. 


Temperature. 

Pressure,  mm. 

Temperature. 

Pressure,  mm. 

F. 

C. 

F. 

C. 

32° 

0° 

4.6 

71.6° 

22° 

19.7 

41 

5 

6.5 

73.4 

23 

20.9 

46.4 

8 

8.0 

75.2 

24 

22.2 

48.2 

9 

8.6 

77.0 

25 

23.6 

50.0 

10 

9.2 

78.8 

26 

25.1 

51.8 

11 

9.8 

80.6 

27 

26.5 

53.6 

12 

10.5 

82.4 

28 

28.1 

55.4 

13 

11.2 

84.2 

29 

29.8 

57.2 

14 

11.9 

86.0 

30 

31.5 

59.0 

15 

12.7 

87.8 

31 

33.4 

60.8 

16 

13.5 

89.6 

32 

35.4 

62.6 

17 

14.4 

91.4 

33 

37.4 

64.4 

18 

15.4 

93.2 

34 

39.6 

66.2 

19 

16.3 

95.0 

35 

41.8 

68  0 

20 

17.4 

69.8 

21 

18.5 

212^6 

100 

760.0 

INDEX 


%*  Acids  are  all  listed  under  "  acid,"  and  salts  under  the  positive  radical. 


ACCUMULATORS,  834 
Acetone,  610 
Acetylene,  592 

torch,  593 
Acid,  acetic,  609,  610,  615 

ionic  equilibrium,  693 
antimonic,  847 
arsenious,  843 
boracic,  638 
boric,  638 
bromic,  486 
carbonic,  575 
Caro's,  449 
chlorauric,  759 
chloroplatinic,  537,  908 
chloric,  482 
chlorous,  483 
cyanic,  626 
dithionic,  449 
formic,  614 
hydrazoic,  521,  541 
hydriodic,  279 
hydrobromic,  274 
hydrochloric,  117,  206 

chemical  properties,  212 
hydrocyanic,  626 
hydrofluoric,  282 

neutralization  of,  396 
hydrofluosilicic,  633 
hypochlorous,  223,  473,  475 

bleaching  by,  477 

neutralization  of,  397 

oxidation  by,  476 

thermochemistry,  479 
hyponitrous,  538 
hypophosphoric,  562 
hypophosphorous,  561 
hyposulphurous,  443 
iodic,  446,  487 


Acid,  metaphosphoric,  557,  560 
metastannic,  827 
muriatic,  207 
nitric,  525 

fuming,  527 

graphic  formula,  540 

nitron  test,  528 

oxidizing  actions,  534 

properties,  526 

test,  530 

nitrosylsulphuric,  433 
nitrous,  537 
orthoarsenic,'  843 
orthophosphoric,  557,  558 
osmic,  905 
oxalic,  616 
pentathionic,  449 
perchloric,  484 
periodic,  488 
permanganic,  881 
peroxidic,  315 
persulphuric,  448 
phosphorous,  561 
picric,  527 

pyrophosphoric,  557,  560 
pyrosulphuric,  437 
salts,  400 

Schiitzenberger's,  444 
selenic,  454 
silicic,  634 
a-stannic,  826 
/8-stannic,  827 
stearic,  619 
sulphuric,  117 

chamber  process,  431 

constitution,  441 

contact  process,  428 

dilute,  properties,  439 

Nordhausen,  437 


911 


912 


INDEX 


Acid,  sulphuric,  properties,  437 

sulphurous,  444 

tannic,  895 

telluric,  455 

tetrathionic,  449 

thiosulphuric,  447 

trithionic,  449 

xanthoproteic,  528 
Acidimetry,  389 
Acids,  117,  324,  373 

action  on  insoluble  sulphides,  419 

action  with  salts,  397 

dibasic,  ionization  of,  439 

chemical  activities  of,  369 

complex,  salts  of,  649 

interaction  with  insoluble  substances, 
713 

non-ionic  formation  of,  406 

organic,  614 

oxygen  acids  of  halogens,  472 
oxidizing  power,  480 

polythionic,  449 

structural  formulae,  562 

sulphur,  431 

weak,  393 
Activity,  chemical,  37 

measured  by  speed  of  reaction,  128 

order,  metals,  129 
non-metals,  284 
Adsorption  of  colloids,  624 
Affinity,    and    double    decomposition, 
383 

electro-measurement  of,  804 

measurement  of,  37 

misuse  of  term,  38,  127 
Air,  a  mixture,  506 

composition  of,  507 

dust  in,  504 

liquid,  509 

saltpeter,  533 

water  vapor  in,  147 
Alabaster,  717 
Alcohol,  ethyl,  607,  608,  616 

methyl,  610,  616 
Alizarin,  820 
Alkalimetry,  387 
Alkaline  earths,  701 
Allotropic  modifications,  315,  551 
Alloys,  644 

anti-friction,  845 


Alumina,  812 
Aluminium,  808 

acetate,  815 

bronze,  739,  809 

chloride,  811 

hydroxide,  811 

-ion,  reactions,  821 

oxide,  812 

sulphate,  813 
Aluminothermy,  653,  810 
Alums,  813 
Alundum,  813 
Amalgams,  644 
Ammonia,  515 

household,  520,  725 

properties,  517 

-soda  process,  688 

synthetic,  516 
Ammonium,  amalgam,  681 

bromide,  678 

carbamate,  680 

carbonates,  679 

chloride,  520,  678 

cyanate,  583,  626 

dichromate,  857,  862 

hydroxide,  519,  520,  678 

iodide,  678 

-ion,  reactions,  681 

molybdate,  863 

nitrate,  679 

oxide,  519 

phosphomolybdate,  561,  864 

sulphantimonate,  848 

sulpharsenate,  844 

sulphate,  680 

sulphides,  680 

sulphostannate,  828 
Amphoteric  hydroxides,  771 
Analysis,  simplified  by  ions,  385 
Anhydride,  giving  several  acids,  487 

mixed,  449,  555 

relation  to  acid  or  salt,  484 
Anhydrite,  717,  878 
Anions,  357 
Anthracene,  613 
Antimony,  845 

halides,  846 

oxides,  847 

salts,  847 

sulphides,  848 


INDEX 


913 


Apatite,  547,  719 

Aqua  regia,  536 

Aqueous  tension,  111 

Aragonite,  704 

Argon,  509 

Arithmetical  problems,  257 

Arsine,  840 

dissociation  of,  296 
Arsenic,  839 

halides,  841 

oxides,  842 

sulphides,  843 
Asbestos,  636 
Asphalt,  587 
Assaying,  761 
Association,  282 
Atmosphere,  499 
Atomic,  numbers,  468 

weight,  of  a  new  element,  261 

weights,  63,  65,  inside  rear  cover 
advantages  of,  245 
determination  of,  239 
history  of,  247 
Atoms,  64 

properties  of,  262 

structure  of,  470 
Attributes,  40 

BABBITT'S  metal,  845 
Baking,  powders,  689 

soda,  689 

Balance,  first  use  of,  10 
Barite,  729 
Barium,  729 

carbonate,  729 

chlorate,  730 

chloride,  730 

peroxide,  82,  316,  731 

sulphate,  730 
Baryta-water,  731 
Bases,  149,  325,  373 

action  with  salts,  398 

chemical  activities  of,  370 

non-ionic  formation  of,  407 

weak,  394 
Basic  oxides,  149 

salts,  401 
Battery,  dichromate,  797 

storage,  834,  see  Cell 
Bauxite,  811 


Beer,  608 

Bell-metal,  739 

Bengal  saltpeter,  525 

Benzene,  613 

Benzine,  586,  593 

Beryl,  636 

Beryllium,  763 

Berzelius,  61 

Bessemer  process,  888 

Birkeland-Eyde  process,  533 

Bismuth,  849 

salts,  850 
Black-lead,  569 
Blast  furnace,  886 
Bleaching,  by  hypochlorous  acid,  477 

not  a  chemical  property,  490 

powder,  475,  716 
Blue  prints,  757 

vitriol,  745 
Body,  definition  of,  7 
Bone-ash,  547 
Boracite,  639 
Borax,  639,  691,  725 
Bordeaux  mixture,  745 
Boron,  637 

halides,  638 

hydrides,  638 
Brandy,  608 
Braunite,  876,  877 
Bricks,  816 
Brin's  process,  82 
Britannia  metal,  824 
Bromine,  preparation,  268 

properties,  270 
Brownian  movement,  622 
Brucite,  766 
B.T.U.,  612 

CADMIUM,  773 

-ion,  reactions,  774 
Caesium,  677 
Calamine,  768 
Calcined  magnesia,  765 
Calcining,  424 
Calcium,  702 

bicarbonate,  576,  705 

carbide,  571,  592,  716 

carbonate,  704 

chloride,  703 

cyanamide,  720 


914 


INDEX 


Calcium,  fluoride,  706 

hydride,  703 

hydroxide,  709 

-ion,  reactions,  727 

nitrate,  717 

nitride,  702 

oxalate,  710 

oxide,  708 

phosphates,  719 

phosphide,  553 

primary  phosphate,  559 

silicate,  725 

sulphate,  717 

sulphide,  719 
Calculations,  74-77 

volumes,  258 

weights,  257 
Caliche,  525 
Calomel,  777,  782 
Calorie,  98 
Calorific  power,  612 
Calorimeter,  98 
Carbides,  570,  655 
Carbon,  566 

chemical  properties,  569 

dioxide,  572 

as  plant  food,  579 
in  air,  501 
properties,  573 
uses,  576,  579 

disulphide,  570 

monoxide,  577 
properties,  578 

oxides  of,  572 

prints,  859 

suboxide,  579 

tetrachloride,  570 
Carbona,  571 
Carbonyl  chloride,  582 
Carborundum,  572 
Carnallite,  664 
Cassiterite,  823 
Castner-Kellner  process,  667 
Catalysis,  97 
Cations,  356 
Cause  in  science,  39 
Celestite,  728 
Cell,  Bunsen,  797 

Clark,  797 

combination,  788,  789 


Cell,  concentration,  789,  801 

Daniell,  796 

displacement,  788,  791,  796 

dry,  797 

Edison,  836 

Leclanche",  797 

oxidation,  788,  792 

potentials  and  concentration,  794 

Weston,  797 

Cells,  facts  concerning  all,  792 
Celluloid,  542 
Cellulose,  603 
Cement,  817 
Cerite,  808 
Cerium,  837 
Chalcocite,  736 
Chalk,  704 
Charcoal,  bone,  610 

wood,  610 
Chemical  activity,  37 

and  iomzation,  369 

methods  of  measuring,  805 
Chemical  affinity,  304 

misuse  of  term,  804 
Chemical  change,  cause  of,  38 

combination,  12,  14 

decomposition,  17 

displacement,  18 

double  decomposition,  20 

energy  in,  28 

internal  rearrangement,  20 

speed  of,  37 

substitution,  225 

varieties  of,  228 
Chemical  equilibrium,  287 

and  concentration,  291 

and  temperature,  304 

characteristics,  289 

displacement  of,  301 

formulation  of,  297 

heterogeneous,  300 

history,  304 

homogeneous,  289 

kinetic  explanation,  289 
Chemical  potential,  803 
Chemical  properties,  inept  statements 
of,  490 

specification  of,  489 
Chemical  relations,  267 

alkaline  earths,  701 


INDEX 


915 


Chenmical  relations,  and  periodic  sys- 
tem, 463 

determine  groups,  464 
Chemically  pure,  6 

Chemistry  and  physics,  distinction  be- 
tween, 48 

Chile  saltpeter,  525 
China,  817 
Chitin,  704 
Chloride  of  lime,  475 
Chlorides,  test  for,  860 

ways  of  making,  213 
Chlorine,  215 

chemical  relations,  226 

dioxide,  483 

monoxide,  473 

nascent,  537 

not  a  bleacher,  478 

preparation,  215-220 

properties,  221-226 
Chloroform,  589 
Chromates,  859 
Chrome-alum,  862 
Chrome-red,  859 
Chrome-tanning,  859 
Chromic  anhydride,  857,  860 

hydroxide,  861 

chloride,  861 

oxide,  861 

sulphate,  862 
Chromite,  855 
Chromium,  854 
Chromous  compounds,  862 
Chromyl  chloride,  860 
Clay,  816 
Coal,  611 
Cobalt,  899 

complex  compounds,  901 
Cobaltic  compounds,  900 
Cobaltous  compounds,  900 
Cochineal,  820 
Coke,  613 
Colemanite,  639 
Collodion,  542 
Colloidal  suspension,  621,  817 

coagulation  of,  622,  844 
Columbium,  852 
Combining  proportions,  52 

measurement  of,  55 
Combining  weights,  57 


Combustion,  91 

spontaneous,  95 
Components,  7 
Composition,  25 
Compound  substances,  22 
Compounds,  molecular,  154,  530 
Complex  acids,  salts  of,  649 
Concurrent  reactions,  483,  485  ' 
Conditions,  40,  67 
Conductivity,  of  electrolytes,  361 
Condy's  fluid,  881 
Congo  red,  392,  820 
Consecutive  reactions,  272,  445,  486 
Constant  boiling-point,  acid  of,  211 
Constant,  equilibrium,  298 

ionization,  359 
Constituents,  12' 
Contact,  action,  97 

agents,  negative,  445 
Copper,  736 

acetylide,  592 

alloys,  738 

ammonio-compounds  of,  739-742 

analytical  reactions,  746 

basic  chloride,  738 

chemical  relations,  734 

complex  cyanides,  744 

electro-refining,  747 

sulphides,  746 
Cordite,  541 
Corn  syrup,  605 
Corrosive  sublimate,  777 
Couples,  119,  800 
Crayon,  704 
Cream  of  tartar,  690 
Critical  temperature,  166 
Crocoisite,  855 
Crops,  fertilization  of,  721 
Cryohydrates,  200 
Crystal  structure,  470 
Crystallization,  water  of,  154 
Crystals,  171-174 
Cupric  acetate,  745 

bromide,  378,  740 

carbonate,  744 

chloride,  739 

hydroxide,  743 

nitrate,  744 

oxide,  56,  743 

sulphate,  745 


916 


INDEX 


Cupric  sulphate,  hydrates,  150-153 

hydrolysis,  399 
Cuprous  chloride,  740 

iodide,  741 

oxide,  743 
Cyanogen,  625,  78 

D ALTON,  61 

Deacon's  process,  217,  305 

Decay,  91 

Decrepitation,  670 

Definitions,    in    experimental    terms, 

238,  264 

Deliquescence,  197 
Density  of  gases,  111 

relative,  260 
Density,  vapor,  113 
Detinning,  824 
Dextrose,  605 
Dialysis,  621 
Diamond,  567 
Diaspore,  811 

Dibasic  acids,  ionization  of,  439 
Diffusion,  163 

gases,  125 

liquids,  168 

pressure,  185 
Dimorphous,  412 
Discharging  potentials,  798 
Disinfectants,  478 
Dissociation,  148 

illustrations,  260 

in  solution,  324,  339 

ionization,  333 
Distillation,  44 

fractional,  587 

in  steam,  563 
Double     decomposition,     formulation, 

380 

Double  salts,  402 
Dust  in  air,  504 
Dyeing,  818 
Dynamite,  541 

EARTHENWARE,  816 
Eau  de  Javd,  669 
Efflorescence,  151 
Electric  furnace,  549 
Electrical  energy,  units,  786 


Electrolysis,  343,  798 
energy  used  in,  359 
explained,  351-355 
of  water,  121 

quantities  of  electricity  used,  357 
Electrolytes,  conductivity  of,  361 
Electrolytic  refining,  799 
Electromotive  chemistry,  786 
Electromotive  series,  metals,  404 
Electrons,  354 

and  oxidation,  493 
Electrophoresis,  622,  817 
Electroplating,  755 
Electrostriction,  397 
Electrotyping,  746 
Element,  22,  23 
Elements,   metallic   and  non-metallic, 

150,  457 

relative  quantities,  24,  114 
Emery,  812 
Emulsion,  624 
Endothermal  changes,  36 
Energy,  30 

and  chemical  change,  28 
chemical  or  internal,  34,  35 
conservation  of,  32 
definitions  of,  32 
factors  of,  803 
free,  35 
units  of,  34 
world's,  580 
Enzymes,  607 
Epsom  salts,  767 
Equation,  70 
writing,  73,  542 

anhydride  method,  496 

partial  equations,  209 

positive    and    negative    valences, 

425,  493 

positive  charges,  496 
Equations,  molecular,  256 
Equilibrium,  169 
chemical,  287 
ionic,  358,  377 

displacement  of,  378 
Equivalent  weights,  63 
Esters,  617 
Ether,  ethyl,  621 

solubility,  180 
methyl,  620 


INDEX 


917 


Ethyl  acetate,  617 

-hydrogen  sulphate,  591,  617 
Ethylene,  590 

bromide,  591 
Evaporation,  15 
Exothermal  changes,  36 
Explanations,  three  kinds,  10 
Extraction,  189 

FACT  in  chemistry,  45 

Fat,  618 

Fehling's  solution,  744 

Felspar,  5,  636 

Fermentation,  607 

Ferrates,  899 

Ferric  compounds,  895 

thiocyanate,  292,  898 
Ferrosilicon,  630 
Ferrous  compounds,  893 

sulphide,  16,  419 
Fertilizers,  673,  720,  721 
Filtration,  15 
Fixation  of  nitrogen,  532 
Flame,  594 

blast  lamp,  597 

Bunsen,  597 

why  non-luminous,  600 

cause  of  luminosity,  599 
Flour,  5 
Fluorine,  280 
Flux,  640,  652 
Foods,  627 

composition  of,  627 

digestion  of,  628 

fuel  values,  628 
Formulae,  calculation  of,  71,  89,  154 

graphic,  322 

reaction,  151 

structural,  322 
Fractions  ionized,  367 

calculation  of,  366 
Franklinite,  768 
Freezing  mixtures,  200 
Freezing-point,  143 

depression  of,  334 
Froth  notation,  737,  768 

GALENA,  829 
Gallium,  807 
Galvanized  iron,  769,  800 


Garnet,  636 
Gas,  blau,  594 

coal,  612 

illuminating,  594 

-lighters,  837 

oil,  594 

Pintsch,  594 

torch,  593 

water,  carburetted,  594 
Gases,  density  of,  111 

ease  of  liquefaction,  427 

heat-capacities  of,  511 

laws  of,  104-110 

liquefaction  of,  507 

mixed,  111 

partial  pressure,  111 

purification,  123 

six  physical  properties,  87 

solubility  in  liquids,  87 

solubilities,  428 
Gasoline,  586 
German  silver,  902 
Germanium,  822 
Glass,  283,  726 

solubility  of,  143 
Glauber's  salt,  690 
Glucinum,  763 
Glucose,  604 
Glycerine,  616 
Glyceryl  nitrate,  528,  540 
G.M.V.,  236 

number  of  molecules  in,  238 
Gold,  758 

compounds  of,  760 

reactions,  761 

Gram-Molecular  Volume,  236 
Granite,  4 
Graphite,  568 
Guano,  525 
Gun-metal,  739 
Guncotton,  528,  617 
Gunpowder,  670 
Gypsum,  717 

'HEMATITE,  885 

Halogens,  267,  284 

chemical  relations,  486 
Hardness,  scale  of,  909 
Hausmannite,  876 
Heavy-spar,  729 


918 


INDEX 


Helium,  511 
Hooke,  8 
Horn-silver,  749 
Humidity,  502 
Hydrargyllite,  811 
Hydrates,  150-154 

conditions  fixing  degree  of  hydration, 
878 

permeable  by  water  vapor,  455 
Hydrazine,  521 
Hydrocarbons,  585 

cracking  of,  593 

paraffins,  585 

saturated,  585 

unsaturated,  590 
Hydrogen,  114 

active  (nascent),  543 

chemical  properties,  126 

dissociation  of  molecules,  253 

-ion,  373 

physical  properties,  124 

preparation,  114-122,  631,  703 
Hydrogen  bromide,  preparation,  271 

properties,  273 

Hydrogen    chloride,     chemical    prop- 
erties, 212 

composition,  227 

physical  properties,  210 

preparation,  206 

preparation,  theory  of,  207-209 
Hydrogen  fluoride,  preparation,  281 
Hydrogen  iodide,  287,  306 

preparation,  277 

properties,  279 
Hydrogen  peroxide,  preparation,  315 

properties,  318 

Hydrogen  pentasulphide,  422 
Hydrogen  sulphide,  414 

properties,  415 

properties  of  solution,  417 
Hydrogenite,  631 
Hydrolysis,  210,  398,  418,  646-648 

aluminium  carbonate,  814 

aluminium  sulphide,  815 
Hydrolyte,  703 
Hydroxide-ion,  374 
Hydroxylamine,  522 
Hypo,  691 
Hypochlorites,  preparation,  474 

properties  of,  480 


Hypotheses,  formulative,  176 
stochastic,  176 

ICELAND  spar,  704 
Impurities,  6 
Indicators,  391 
Indigo,  819 

oxidation  of,  314,  476,  479,  538 
Indium,  807 

Infusible  white  precipitate,  781 
Infusorial  earth,  633 
Ink,  India,  597 

marking,  754 

printer's,  597 

writing,  895 
lodimetry,  277 
lodic  anhydride,  487 
Iodine,  chlorides  of,  285 

dissociation,  276 

preparation,  274 

properties,  275 

titration  of,  448 

vapor,  261 
lodoform,  589 
Ionic  equilibrium,  358,  377 

displacement  of.  378 

excess  of  one  ion,  694 

saturated  solutions,  697 

viewed  quantitatively,  693 
Ionic,  migration,  345 

reactions,  kinds  of,  402 

substances,  chemical  properties,  384 

nomenclature,  356 
lonization,  342 

and  chemical  activity,  369 

degree  of,  366 

repression  of,  694 
lonogens,  355 

classes  of,  372 

non-ionic  formation  of,  406 
Ions,  are  substances,  375 

displacement  of,  403 

nature  of,  348 
Iridium,  906 
Iron,  884 

carbonyls,  898 

cast,  887 

chemical  properties,  892 

magnetic  oxide,  89,  116,  896 

metallurgy,  885 


INDEX 


919 


Iron,  passive,  893 

reactions,  899 

-stone,  885 

wrought,  887 
Isomers,  583 
Isomorphism,  657 
Isoprene,  590 

KAINITE,  673 

Kaolin,  636,  816 

Kelp,  664 

Kieserite,  767 

Kindling  conditions,  95 

Kinetic-molecular  theory,  160 

Krypton,  511 

LAKES,  820 
Lampblack,  596 
Laughing  gas,  540 
Lavoisier,  8,  22,  80 
Law,  Avogadro's,  163,  231 

Boyle's,  104,  107,  161 
deviations  from,  164 

Charles',  110,  162 

combining  weights,  59 

component  substances,  6 

conservation  of  mass,  52 

constant  heat  summation,  100 

definite  proportions,  54 
exception  to,  54,  871,  872 

Dulong  and  Petit's,  245 

Faraday's,  350 

Gay-Lussac's,  157,  163 

Henry's,  188 

in  chemistry,  45 

ion-product  constancy,  698 

Le  Chatelier's,  307 

mass  action,  292 

mobile  equilibrium,  306 

molecular  concentration,  292 

multiple  proportions,  58 

of  octaves,  459 

of  partition,  189 

periodic,  462 

phase  rule,  705 

specific  physical  properties,  4 

thermoneutrality,  381 

transformation  by  steps,  544 

van't  Hoff's,  305 
Laws,  two  kinds,  232 


Lead,  829 

acetate,  834 

carbonate,  833 

halides,  830 

-ion,  reactions,  836 

nitrate,  833 

radio-,  871,  872 

oxides,  831 

sugar  of,  834 

sulphate,  834 

sulphide,  834 
Le  Blanc  process,  686 
Light,  and  chemical  change,  581 

chemical  effect  of,  581,  757 
Lignin,  603 
Limestone,  705 
Limewater,  709 
Liquefaction  of  gases,  507 
Liquid  air,  509 
Liquid  and  vapor,  168 
Lithium,  692 
Lithopone,  730 
Litmus,  392 

Lomonossov,  6,  8,  22,  167 
Lunar  caustic,  753 

MAGNALIUM,  809 
Magnesia  alba,  766 
Magnesium,  766 

carbonates,  766 

chloride,  765 

hydroxide,  766 

interaction  with  ammonium  salts, 
766 

-ion,  reactions,  768 

nitride,  514 

oxide,  765 

phosphates,  767 

sulphate,  767 

sulphide,  767 
Magnetite,  885 
Malachite,  736 
Manganates,  880 
Manganese,  875 

bronze,  739 

heptoxide,  882 
Manganic  compounds,  879 
Manganin,  876 
Manganites,  879 
Manganous  compounds,  877 


920 


INDEX 


Manganous  dithionate,  450 

Marsh's  test,  841,  846 

Matches,  552 

Mayow,  8 

Mendelejeff's   Table,    461,   opp.    rear 

cover 
Mercuric  chloride,  777 

cyanide,  781 

fulminate,  541,  781 

oxide,  17 

thiocyanate,  781 
Mercurous  chloride,  777 
Mercury,  775 

ammono-compounds,  781 

iodides,  779 

nitrates,  780 

oxides,  779 

reactions,  782 

sulphides,  780 
Metallic  elements,  150,  457 

chemical  relations,  645 

classes  of,  650 

of  alkalies,  chemical  relations,  661 
Metals,  compounds  of,  hydroxides,  654 

oxides,.  653 
Metals,  electric  conductivity,  644 

extraction  from  ores,  652 

melting-points,  643 

occurrence,  652 

of  the  alkalies,  661 

physical  properties,  642 

salts  of,  654 

world's  production,  645 
Metastable  condition,  194 
Methane,  588 

synthesis,  570 
Methods,  experimental,  42 
Methyl  formate,  617 

orange,  392 
Mica,  4,  636 

Microcosmic  salt,  559,  681 
Migration,  ionic,  345 

velocities,  347 
Mirrors,  silvering  of,  755 
Mispickel,  839 
Minium,  832 
Mixtures,  6 
Molecular,  compounds,  154,  530 

equations,  256 

formulae,  249-253 


Molecular,  theory,  160 

gases,  160-167 

gases,  summary,  165 

history,  167 

liquids,  168 

solids,  170 

solutions,  183 
weights,  231 

freezing-point  method,  336 
Molecules,  64 
Molybdenum,  863 
Monde  process,  903 
Monel  metal,  902 
Mordants,  820 
Mortar,  700 

NAPHTHA,  586 

Naphthalene,  613 

Nascent  (active),  hydrogen,  543 

chlorine,  537 

oxygen,  480 
Natural  gas,  588 
Negative  elements,  457 
Neon,  511 

Nessler's  reagent,  782 
Neutralization,  in  conductivity  trough, 
388 

ionic  formulation,  386 

of  little  ionized  substances,  394 

thermal  effects,  388,  396 

volume  change  in,  397 
Nickel,  901 
Niobium,  852 
Niton,  511,  871 
Nitric  anhydride,  527 

oxide,  529,  532 
Nitrides,  514 
Nitro-derivatives,  527 
Nitro-lime,  720 
Nitrogen,  active  form,  515 

chemical  relations,  513 

fixation  of,  515,  532 

iodide,  523,  541 

preparation,  513 

properties,  514 

tetroxide,  531 

trichloride,  523 
Nitroglycerine,  528,  540,  617 
Nitron,  528 
Nitrosyl  chloride,  537 


INDEX 


921 


Nitrous,  anhydride,  538 

oxide,  539 

Nomenclature,  ionic  substances,  356 
Non-metallic  elements,  150,  457 

OILS,  drying,  620 

vegetable,  618 
Oleum,  437 

Open-hearth  process,  889 
Orpiment,  839 
Osmium,  905 
Osmotic  pressure,  327-334 

approximate  relations,  332 
exact  relations,  331 
measurement  of,  330 

suction,  329 
Oxidation,  91 

accompanies  reduction,  417 

and  reduction,  491-493 
Oxides,  acidic,  150 

basic,  149 

nomenclature,  90 

Oxidizing  agents,  why  active,  314, 
Oxone,  85 
Oxygen,  79 

chemical  properties,  87 

history  of,  79 

in  respiration,  92 

"nascent,"  224,  480 

physical  properties,  86 

preparation,  81-85 

substances  indifferent  to,  92 

uses,  92 
Ozocerite,  587 
Ozone,  311,  315 

PAINT,  730,  836 

luminous,  719 
Palladium,  906 
Paper,  603,  815 
Parke's  process,  749 
Pattinson's  process,  749 
Pauling  process,  534 
Pearl  ash,  672 
Perchloric  anhydride,  484 
Periodic  system,  456 

grouped  by  relations,  464 

inadequacy  of,  467 

properties  and  relations  in,  462 

uses,  465 


Permanganates,  881 

Permutite,  724 

Peroxides,  constitution  of,  321 

Petrol,  586 

Petrolatum,  587 

Petroleum,  586 

Pewter,  824 

Phase  rule,  705 

Phases,  192 

Phenol,  613 

Phenolphthalem,  391 

Phlogiston,  9 

Phosgene,  225 

Phosphine,  552 

Phosphonium  compounds,  554 

Phosphorite,  719 

Phosphorus,  chemical  relations,  547 

oxychloride,  555 

pentabromide,  555 

pentachloride,  260,  299,  554 

pentoxide,  556 

preparation,  547 

red,  properties,  549,  551 

sulphides,  563 

trichloride,  554 

trioxide,  556 

white,  properties,  549,  550 
Photochemistry,  581 
Photography,  755 

blue  prints,  898 

Physics  and  chemistry,  distinction  be- 
tween, 48 

Physics  in  chemistry,  47,  84 
Plaster  of  Paris,  718 
Platinum,  907 

prints,  757 

Polarization,  360,  798 
Polymerization,  282 
Polymorphous,  412 
Polysulphides,  421 
Porcelain,  816 
Positive  elements,  457 
Potash,  671 
Potassium,  662 

bromate,  669 

bromide,  665 

carbonate,  671 

chlorate,  83,  480,  669 

chloride,  663 

chromate,  856 


922 


INDEX 


Potassium,  cobalticyanide,  901 

cobaltinitrite,  901 

cobaltocyanide,  901 

cuprocyanide,  744 

cyanate,  626,  673 

cyanide,  672 

dichromate,  856 

ferricyanide,  897 

ferrocyanide,  897 

fluoride,  665 

-hydrogen  tartrate,  675,  690 

hydroxide,  665,  700 

hypobromite,  486 

hypochlorite,  474,  669 

iodate,  669 

iodide,  664 

-ion,  reactions,  675 

manganate,  880 

nitrate,  525,  670 

oxides,  668 

percarbonate,  672 

perchlorate,  483 

permanganate,  881 
as  oxidizer,  882 

pyroantimoniate,  848 

pyrosulphate,  674 

sulphates,  673 

sulphides,  674 

thiocyanate,  626,  673 

zincate,  770 
Potential  differences,  metals,  793 

non-metals,  797 

Precipitates,  description  of,  214 
Precipitation,  formulation  of,  382 

of  ionogens,  382 

rule  for,  711 

theory  of,  710 
Pressure,  osmotic,  327-334 
Principles,  summary  of,  229,  309 
Priestley,  80 
Producer  gas,  577 
Properties,  specific  physical,  3 

list  of  specific  physical,  40 
Proteins,  515,  520,  528 
Proustite,  749 
Prussian  blue,  897 
Purple  of  Cassius,  825 
Pyrargyrite,  749 
Pyrene,  571 
Pyrolusite,  876 


Pyrosulphates,  438 

QUALITATIVE  analysis,   recognition   of 

cations,  783 
Quartation,  761 
Quartz,  4,  633 
Quartz  glass,  634 
Quicklime,  708 
Quicksilver,  776 

RADICALS,  117 

organic,  589 
Radium,  865 

emanation,  512 
Reaction,  formulae,  151 

speed,  320 

and  temperature,  304 
lect.  exp.,  446 
Reactions,  concurrent,  483,  485 

consecutive,  445,  486 

ionic,  kinds  of,  402 
Realgar,  839,  844 
Red  lead,  832 

Reduction  and  oxidation,  491-493 
Refrigeration,  517 
Rey,  Jean,  8 
Rhodium,  906 
Rinmann's  green,  771 
Roasting,  424 
Rochelle  salt,  690 
Root  nodules,  515 
Rubber,  artificial,  590 
Rubidium,  677 
Ruby,  812 
Rusting,  801 

explanation  of,  8 

of  metals,  7,  9 
Ruthenium,  905 

SALERATUS,  672 

Salt,  inept  definition  of,  407 

Salts,  214,  324,  375 

acid,  400 

action  with  acids  and  bases,  397 

action  with  water,  398 

basic,  401 

chemical  activity  of,  369 

common  hydrates,  657 

double,  402 

mixed,  401 


INDEX 


923 


Salts,  non-ionic  formation  of,  407 

Sanitation,  bleaching  powder  in,  478 

JSaponification,  618 

Sapphire,  812 

Scandium,  807 

Scheele,  80 

Schlippe's  salt,  848 

Schoenite,  673 

Schwerin  process,  817 

Scientific  method,  3 

Selenite,  717 

Selenium,  453 

compounds,  454 
Serpentine,  636 
Sewage  disposal,  92 
Sherardizing,  760 
Siderite,  885 

Siemens-Martin  process,  889 
Silicon,  630 

-bronze,  739 

dioxide,  633 

hydride,  631 

properties,  631 

tetrachloride,  632 

tetrafluoride,  632 
Silk,  artificial,  542,  743 
Silver,  748 

alum,  754 

carbonate,  754 

chemical  relations,  748 

colloidal,  750 

complex  salts,  752 

halides,  751 

-ion,  reactions,  757 

nitrate,  18,  753 

oxide,  753 

peroxidate,  754 

plating,  755 

sulphate,  754 

sulphide,  754 
Simple  substances,  22 

preparation,  122 
Stability,  148 
Smokeless  powder,  541 
Soap,  chemical  properties,  620 

cleansing  power,  623 

colloidal,  623 

hydrolysis  of,  399 
Soapstone,  636 
Soda  water,  574 


Sodium,  683 

aluminate,  812 

amalgam,  684,  776 

bicarbonate,  575,  689 

carbonate,  686 

chloride,  18,  684 

cyanide,  721 

dichromate,  857 

hydride,  443,  684 

hydroxide,  685 

hypochlorite,  474 

hyposulphite,  443 

-ion,  reactions,  692 

metaphosphate,  559,  560 

nitrate,  685 

nitrite,  537,  685 

orthophosphates,  558 

oxide,  685 

periodate,  488 

peroxide,  85,  685 

phosphates,  691 

plumbite,  832 

pyrophosphate,  559 

silicates,  634,  691 

a-stannate,  826 

/3-stannate,  827 

sulphate,  690 

solubilities,  193,  195 

sulphide,  hydrolysis,  418 

tetraborate,  639,  691 

tetrathionate,  448 

thiosulphate,  447,  691 
Solder,  830 

Solubilities,  bases  and  salts,  656,  in- 
side front  cover 

gases,  187,  428 

gases  in  liquids,  87 

liquids  in  water,  180 

salts,  191,  656,  inside  front  cover 
Solubility,  independent,  187 

insoluble  salts,  802 

measurement  of,  180 

-product  constant,  698 

separation  by,  481 

temperature  and,  190 

units  used,  181 
Solution,  178 

boiling-points  of,  198,  337 

concentration.  181 

definition,  203 


924 


INDEX 


Solution,  density  of,  201 

freezing-point  of,  199,  334 

heat  of,  202 

heat  of,  data,  203 

insoluble  salts,  to  complex  ions,  741 

molar,  183 

normal,  182 

of  insoluble  substances,  712 

osmotic  pressure  of,  327 

physical  or  chemical,  202 

pressure,  186 

rule  for,  711 

saturated,  181,  192 
definitions  of,  194 

suction,  186 

supersaturated,  192 

vapor  tension  of,  197 

varieties  of,  179 

volume  changes,  201 
Solvay  process,  688 
Solvents,  ionizing,  357 
Spectroscope,  675 
Specularite,  885 

Speed  of  reaction,  and  concentration, 
96  * 

and  temperature,  93 
Spinells,  812 
Spinthariscope,  867 
Stability,  93 
Stalactites,  705 
Standard  solutions,  390 
Stannic  chloride,  825 
Stannous  chloride,  825 
Starch,  604 

States  of  aggregation,  48 
Stationary  layer,  503,  597 
Stearin,  619 
Steel,  alloys,  892 

crucible,  890 

tempering,  890 
Stereotype  metal,  747 
Stibine,  846 
Stibnite,  845 
Storage  battery,  834 
Strontium,  728 
Sublimation,  275,  555 
Substance,  definition,  4 

incorrect  definitions,  25 

simple  or  compound,  22 
Sucrose,  605 


Sugar,  cane,  605 

grape,  605 

invert,  606 
Sulphates,  441 
Sulphides,  418 

acids  on  insoluble,  419 

insoluble,  classification,  421 

of  metals,  solubilities  of,  774 
Sulphites,  446 
Sulphur,  amorphous,  413 

chemical  properties,  413 

chemical  relations,  422 

chlorides,  450 

monoclinic,  411 

dioxide,  preparation,  424 
properties,  426 

family,  chemical  relations,  456 

manufacture,  410 

oxidation  by  air,  446 

oxides  of,  424 

physical  properties,  411 

rhombic,  411 

trioxide,  preparation,  428 
properties,  430 

vapor,  261 

viscous,  412 

Sulphuryl  chloride,  427,  451 
Superphosphate,  559,  720 
Sylvanite,  758 
Sylvite,  664 
Symbols,  69 
Synthesis,  566 

TALC,  636 

Tantalum,  852 
Tellurium,  455 
Temperature,  and  reaction  speed,  93 

critical,  166 

Temperatures,  °F.  and  °C.,  910 
Thallium,  807 
Thermite,  810 
Thermochemistry,  98-101,  321 

of  neutralization,  396 
Thionyl  chloride,  450 
Thorium,  837,  872 
Tin,  823 

concentration  cell,  801 

oxides,  827 

-plate,  800 

reactions,  828 


INDEX 


925 


Tin,  sulphides,  828 
Titanium,  837 
Titration,  390 
T.N.T.,  527 
Toluene,  590 
Touch-paper,  671 
Transformation  by  steps,  544 
Triads,  458 
Trinitrotoluene,  527 
Tripoli  powder,  633 
Tungsten,  864 
TurnbulPs  blue,  898 

ULTRAMARINE,  818 

Ultramicroscope,  621 

Units,  molecular  concentration,  294 

weight  and  volume,  74,  909 
Unstable,  misuse  of  term,  148 
Uranium,  865 
Urea,  582 

VALENCE,  130-139 
inept  definitions,  138 
oddities  connected  with,  137 

Van  der  Waals,  165 

Vanadium,  852 

Vapor  pressure,  145 

Vapor  tension,  145 

Varieties  of  chemical  change,   12,  17, 
18,  20,  21 

Vaseline,  587 

Ventilation,  502 

Vermilion,  780 

Vinegar,  609 

Vitriols,  771 

Volumetric  methods,  391 

WASHING  soda,  688,  724 
Water,  141 

action  on  salts,  398 

associated  liquid,  202 

chemical  properties,  147-151 

composition  of,  155-157 


Water,  electrolysis  of,  798 
fraction  ionized,  368 
gas,  577 

carburetted,  594 
glass,  634 
hard,  722 

damage  due  to,  722 
in  laundry,  723,  724 
treatment  of,  723 
of  crystallization,  154 
physical  properties,  143 
purified  by  coagulation,  8 
vapor  pressures  of,  910 

Waterproofing,  816 

Waters,  natural,  141 
purification  of,  142 

Weights,  and  measures,  909 
atomic,  63,  65 
combining,  57 
equivalent,  63 

Weldon  process,  879 

White  lead,  833 

Whiskey,  608 

Wine,  607 

Wireless  electric  waves,  469 

Witherite,  729 

Wood,  distillation  of,  610 

Wood's  metal,  849 

X-RAYS,  469 
Xenon,  511 

ZINC,  768 

blende,  768 

carbonates,  771 

chloride,  769 

hydroxide,  770 

-ion,  reactions,  772 

oxide,  770 

sulphate,  771 

sulphide,  772 
Zincite,  768 
Zircon,  636 
Zirconium,  837 


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